Test Bank for General Chemistry, Reactions First, 1st Edition By Kevin Revell (Macmillan) (All Chapt

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Chapter 22 Chemical Analysis 1. Which piece of glassware would you use for precise volume measurements? a. beaker b. Erlenmeyer flask c. volumetric pipet d. dropper bottle e. round bottom flask ANSWER: c 2. A calibration curve is used in chemical analysis to a. determine the precise volume of the analyte. b. convert the measured signal into the corresponding analyte concentration. c. identify the different components present in a mixture. d. calculate the detection limit of the method. e. measure the absorbance of a solution at different wavelengths.

ANSWER: b 3. How would you report the volume of liquid dispensed by a 10-mL volumetric pipet? a. 10 mL b. 10.0 mL c. 10.00 mL d. 10.000 mL e. 10.0000 mL ANSWER: c 4. In chemical analysis, what is an analyte? a. The solution of known concentration used for comparison. b. The instrument used to measure the signal. c. The specific substance or component being measured or detected. d. The graph showing the relationship between signal and concentration. e. The process of preparing a sample for analysis.

ANSWER: c 5. The detection limit of an analytical method is best defined as a. the maximum amount of an analyte that can be accurately measured. b. the minimum amount of an analyte that can be reliably distinguished from the blank. c. the concentration of an analyte that produces a signal equal to the noise level. d. the range of analyte concentrations over which the method is accurate. e. the point at which the signal becomes saturated and no longer increases with Copyright Macmillan Learning. Powered by Cognero.

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Chapter 22 Chemical Analysis concentration.

ANSWER: b 6. What does the saturation limit refer to in spectroscopy? a. The point at which the sample becomes cloudy and light scattering occurs. b. The minimum concentration of analyte that can be detected. c. The wavelength at which the analyte absorbs the most light. d. The maximum volume of solution that can be held in the cuvette. e. The point at which the absorbance no longer increases linearly with concentration.

ANSWER: e 7. The sensitivity region of an analytical method is characterized by a. a high degree of linearity between the signal and the analyte concentration. b. a low signal-to-noise ratio. c. a narrow range of analyte concentrations that can be measured. d. a high detection limit. e. a significant degree of interference from other substances.

ANSWER: a 8. A chemist prepares a solution by dissolving 5.00 g of potassium permanganate (KMnO4) in enough water to make 250.0 mL of solution. Next, they transferred 2.00-, 4.00-, 8.00-, and 10.00-mL aliquots of this solution to four 50.00-mL volumetric flasks and diluted to the mark with water. What is the concentration of the solution prepared with 2.00 mL of the stock solution? a. 0.126 M b. 0.00506 M c. 0.0101 M d. 0.00253 M e. 0.0202 M ANSWER: b 9. A biochemist prepares a 12.5-mM solution of potassium phosphate. They then transfer 1.00-, 2.00-, 4.00-, and 6.00-mL aliquots from this solution to four separate 10.0-mL volumetric flasks, diluting each to the mark with water. What is the concentration of the solutions prepared with 4.00 mL of the stock solution? a. 1.25 mM b. 2.50 mM c. 5.00 mM d. 7.50 mM e. 10.0 mM Copyright Macmillan Learning. Powered by Cognero.

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Chapter 22 Chemical Analysis ANSWER: c 10. Which of the following R2 values would signify that the data points match up with the function of the calibration curve the most? a. R2 = 0.9987 b. R2 = 0.9925 c. R2 = 0.5000 d. R2 = 0.7824 e. R2 = 0.0012 ANSWER: a 11. Which of the following R2 values would signify an inaccurate calibration curve? a. R2 = 0.9912 b. R2 = 0.9937 c. R2 = 0.9931 d. R2 = 0.9401 e. R2 = 1.0000 ANSWER: d 12. What instrument can measure the mass of chemicals accurately, usually to the nearest 0.1 mg?

ANSWER: electronic balance 13. A solution is prepared by dissolving 7.50 g of ammonium nitrate (NH4NO3) in enough water to make 200.0 mL of solution. A chemist then takes 1.00-, 3.00-, 5.00-, and 7.00-mL aliquots from this solution and transfers them to four different 100.0-mL volumetric flasks, diluting each to the mark with water. Determine the concentration of each resulting solution in mM. ANSWER: Stock Volume (mL) Concentration (mM) 1.00 4.69 3.00 14.1 5.00 23.43 7.00 32.8 14. A chemist prepares a 0.500-M stock solution of glucose. They need to produce four standard solutions with concentrations of 0.100 M, 0.200 M, 0.300 M, and 0.400 M, each with a final volume of 25.0 mL. What volume of the stock solution should be used to prepare each of the standard solutions?

ANSWER: Concentration (M) Volume (mL) 0.100 5.00 0.200 10.0 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 22 Chemical Analysis 0.300 0.400

15.0 20.0

15. A biochemist prepares a 0.500-μM citric acid solution and transfers 2.50-, 5.00-, 7.50-, and 10.00-mL aliquots into four separate 20.00-mL volumetric flasks, filling each flask to the mark with water. What are the final concentrations of citric acid in each solution?

ANSWER: Stock Volume (mL) Concentration (μM) 2.50 0.0625 5.00 0.125 7.50 0.188 10.0 0.250

16. Given the measurements 2.45 g, 2.50 g, 2.55 g, 2.40 g, and 2.60 g, what is the mean of these measurements? a. 2.50 g b. 2.45 g c. 2.55 g d. 2.48 g e. 2.52 g ANSWER: a 17. If the concentrations of a solution were determined to be 1.5 μM, 1.7 μM, 1.6 μM, and 1.4 μM, what is the standard deviation of this data set? a. 0.13 μM b. 0.36 μM c. 0.079 μM d. 0.20 μM e. 0.14 μM ANSWER: a 18. What is the number of degrees of freedom for a data set with three individual measurements? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: b 19. What is the number of degrees of freedom for a data set with 15 individual measurements?

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Chapter 22 Chemical Analysis a. 1 b. 14 c. 15 d. 16 e. 30 ANSWER: b 20. What type of errors limit the precision of measurements? a. random errors b. theoretical errors c. conceptual errors d. calculation errors e. computational errors ANSWER: a 21. What is the mean of the following data set?

10.0, 12.0, 9.50, 14.0, 15.0 a. 9.50 b. 15.0 c. 12.1 d. 11.7 e. 14.0 ANSWER: c 22. What is the standard deviation of the following data set?

10.1, 12.3, 11.7, 12.0, 13.1 a. 10.1 b. 12.3 c. 11.8 d. 1.10 e. 0.875 ANSWER: d 23. What does the variable n stand for in the equations for the mean and standard deviation? a. number of individual measurements b. the mean c. the standard deviation d. the Student's t-test value e. number of degrees of freedom ANSWER: a Copyright Macmillan Learning. Powered by Cognero.

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Chapter 22 Chemical Analysis 24. After conducting 15 trials, your lab finds the concentration of lead in a water sample to be 12.54 ppb, with a standard deviation of 0.15. Report the range of concentrations with a 95% confidence interval. a. 12.54 ± 0.08 ppb b. 12.54 ± 0.06 ppb c. 12.54 ± 0.02 ppb d. 12.12 ± 0.08 ppb e. 12.06 ± 0.04 ppb ANSWER: a 25. After conducting 20 trials, your lab finds the concentration of mercury in a fish sample to be 0.89 ppm, with a standard deviation of 0.08. Report the range of concentrations with a 90% confidence interval. a. 0.84 to 0.92 ppm b. 0.86 to 0.92 ppm c. 0.86 to 0.94 ppm d. 0.86 to 0.96 ppm e. 0.84 to 0.94 ppm ANSWER: b 26. A student performs an experiment to determine the concentration of copper ions in a solution, conducting 15 trials. The average concentration is found to be 1.25 mg/L, with a standard deviation of 0.15 mg/L. Calculate the range of concentrations with a 95% confidence interval.

ANSWER: 1.17 to 1.33 mg/L 27. A research team measures the levels of dissolved oxygen in a lake at various depths. After 12 measurements, they find an average dissolved oxygen concentration of 8.7 ppm, with a standard deviation of 0.6 ppm. Report range of concentrations with a 99% confidence interval.

ANSWER: 8.16 to 9.24 ppm 28. A student performs an experiment to determine the concentration of copper ions in a solution, conducting 15 trials. The average concentration is found to be 1.25 mg/L, with a standard deviation of 0.15 mg/L. Calculate the range of concentrations with a 95% confidence interval.

ANSWER: 1.17 to 1.33 mg/L 29. What does the standard deviation refer to in chemical analysis?

ANSWER Standard deviation is a measure that describes how widely measured values : deviate from the mean value.

30. What does the confidence interval refer to in chemical analysis?

ANSWER The confidence interval is the statistical measure that conveys the likelihood : that the true mean of a series of measurements lies within a certain range of Copyright Macmillan Learning. Powered by Cognero.

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Chapter 22 Chemical Analysis values.

31. Which of the following statements best describes the fundamental principle underlying chromatographic separation? a. Compounds are separated based on their differences in density. b.Compounds are separated based on their differences in solubility in a given solvent. c. Compounds are separated based on their differences in boiling point. d.Compounds are separated based on their differences in affinity for a stationary phase and a mobile phase. e. Compounds are separated based on their differences in reactivity with a specific reagent.

ANSWER: d 32. In a typical chromatographic experiment, the component that has the shortest retention time is a. the component with the highest affinity for the stationary phase. b. the component with the lowest affinity for the stationary phase. c. the component with the highest molar mass. d. the component with the lowest molar mass. e. the component with the highest boiling point.

ANSWER: b 33. In a typical chromatographic experiment, the component that has the longest retention time is a. the component with the highest affinity for the stationary phase. b. the component with the lowest affinity for the stationary phase. c. the component with the highest molar mass. d. the component with the lowest molar mass. e. the component with the highest boiling point.

ANSWER: a 34. In gas chromatography, the mobile phase is typically a. a liquid solvent. b. a solid adsorbent. c. an inert gas. d. a vacuum. e. a supercritical fluid.

ANSWER: c 35. In high-pressure liquid chromatography (HPLC), the mobile phase is typically a. a liquid solvent. b. a solid adsorbent.

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Chapter 22 Chemical Analysis c. an inert gas. d. a vacuum. e. a supercritical fluid.

ANSWER: a 36. What type of instrumental technique uses high pressure to separate substances based on their polarity? a. GC b. UV-vis c. HPLC d. MS e. NMR ANSWER: c 37. Which of the following molecules will have the shortest retention time in a GC experiment? a. C5H12 b. C6H14 c. C7H16 d. C8H18 e. C9H20 ANSWER: a 38. Which of the following molecules will have the longest retention time in a GC experiment? a. C5H12 b. C6H14 c. C7H16 d. C8H18 e. C9H20 ANSWER: e 39. In HPLC, the stationary phase is typically a. a liquid solvent. b. a solid adsorbent. c. an inert gas. d. a vacuum. e. a supercritical fluid.

ANSWER: b 40. What is the size of the detector signal’s peak proportional to in a chromatography experiment?

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Chapter 22 Chemical Analysis a. retention time b. attraction to the mobile phase c. attraction to the stationary phase d. concentration of the analyte e. concentration of the mobile phase ANSWER: d 41. How are molecules separated in a gas chromatography experiment?

ANSWER Molecules are separated based on their volatility and interactions with the : stationary phase in the column.

42. How are molecules separated in a liquid chromatography experiment?

ANSWER In liquid chromatography, molecules are separated based on their polarity, : solubility, and interactions with both the mobile phase and the stationary phase.

43. The following data was recorded for a gas chromatography experiment. Rank the molecules in order of their concentration.

Molecule Retention time (min) A 12.2 B 4.7 C 1.3 ANSWER: C < A < B

Detector signal 86 142 27

44. The following data was recorded for a HPLC experiment. Rank the molecules in order of their attraction toward the mobile phase, from weakest attraction to strongest attractions.

Molecule Retention time (min) W 4.8 X 7.4 Y 3.2 Z 9.2 ANSWER: Y < W < X < Z

Detector signal 176 142 89 199

45. What does the acronym HPLC stand for?

ANSWER: high pressure liquid chromatography 46. What is the absorbance for a sample that had a measured transmission of 0.72 by a UV-vis spectrophotometer? a. 0.16 b. 0.14 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 22 Chemical Analysis c. 0.98 d. 1.7 e. 1.3 ANSWER: b 47. What is the fraction of the incident light transmitted through the sample? a. absorbance b. λmax c. transmittance d. fluorescence e. molar absorptivity ANSWER: c 48. What is the percent transmission for a sample that had a measured absorbance of 1.12 by a UV-vis spectrophotometer? a. 19 % b. 12 % c. 47 % d. 7.6 % e. 9.2 % ANSWER: d 49. A solution of a compound is analyzed in a 1.00-cm cuvette using a UV-vis spectrophotometer. The absorbance of the solution at a wavelength of 520 nm is found to be 0.450. If the molar absorptivity of the compound at 520 nm is 2.50×103 cm–1 M–1, what is the concentration of the compound in the solution? a. 2.27×10–4 M b. 1.80×10–4 M c. 3.40×10–4 M d. 1.52×10–4 M e. 4.59×10–4 M ANSWER: b 50. A student prepares a solution of a colored dye and measures its absorbance using a UV-vis spectrometer.

The student then dilutes the solution by a factor of 2 and measures the absorbance of the diluted solution. How will the absorbance of the diluted solution compare to the absorbance of the original solution? a. The absorbance will be four times as large. b. The absorbance will be twice as large. c. The absorbance will be half as large.

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Chapter 22 Chemical Analysis d. The absorbance will be one-fourth as large. e. The absorbance will remain the same.

ANSWER: b 51. A chemist is analyzing a newly synthesized organic compound. A solution of the compound with a concentration of 1.25×10–4 M has an absorbance of 0.350 at a wavelength of 400 nm in a 1.00-cm cuvette.

What is the molar absorptivity of the compound at this wavelength? a. 2.80×103 cm–1 M–1 b. 1.80×103 cm–1 M–1 c. 3.47×103 cm–1 M–1 d. 7.20×103 cm–1 M–1 e. 8.09×103 cm–1 M–1 ANSWER: a 52. Why is atomic absorption spectroscopy highly selective for the element being analyzed? a. Because the hollow-cathode lamp emits light at all wavelengths. b. Because the flame atomizes all elements in the sample. c. Because the sample is diluted before analysis. d. Because each element has a unique set of electronic energy levels. e. Because the detector is only sensitive to the analyte.

ANSWER: d 53. Which of the following statements about fluorescence is true? a. The emitted photon has a higher energy than the absorbed photon. b. The emitted photon has the same energy as the absorbed photon. c. The emitted photon has a lower energy than the absorbed photon. d. Fluorescence involves the absorption of light in the microwave region. e. Fluorescence is a type of scattering.

ANSWER: c 54. What is the relationship between the excitation and emission wavelengths in fluorescence spectroscopy? a. The excitation wavelength is directly proportional to the emission wavelength. b. The excitation wavelength is longer than the emission wavelength. c. The excitation wavelength is shorter than the emission wavelength. d. The excitation and emission wavelengths are always equal. e. The excitation wavelength is longer at low concentrations and shorter at high concentrations.

ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 22 Chemical Analysis 55. What is the function of the hollow-cathode lamp in atomic absorption spectroscopy? a. To provide a continuous spectrum of electromagnetic radiation. b. To emit light at specific wavelengths absorbed by the analyte. c. To atomize the sample. d. To measure the amount of light transmitted through the sample. e. To produce the excitation wavelength needed for fluorescence.

ANSWER: b 56. A 0.00750-M solution shows an absorbance of 1.12 at 450 nm when measured in a 1.00-cm-path-length cuvette. Determine the percent transmittance of this solution and the molar absorptivity of the solute.

ANSWER: %T = 7.59 % ε = 149 cm–1 M–1 57. What is the absorbance, transmittance, and percent transmittance when the path length is 1.00 cm for a 50.0mM solution with a molar absorptivity (ε) of 18.0 cm–1 M–1 at a 475 nm?

ANSWER: A = 0.900 T = 0.126 %T = 12.6% 58. What is the concentration of a solution that contains an analyte with a molar absorptivity of 25.5 cm–1 mM–1 at 575 nm? The absorbance of the solution was measured to be 1.243 using a spectrometer with a path length of 2.50 cm.

ANSWER: 0.0195 mM 59. According to Beer's law, what are the relationships between absorbance and molar absorptivity, path length, and concentration?

ANSWER Absorbance is directly proportional to molar absorptivity, path length, and : concentration. As the values of those factors increase, so does absorption.

60. A calibration curve was prepared using data from a UV-vis spectroscopy experiment where the absorption was measured for several standard solutions with concentrations ranging from 10.0 μM to 50.0 μM. The spectrometer was equipped with 1.00-cm-wide cuvette. The curve produced the equation y = 123.4x + 0.000123. What is the molar absorptivity of the molecule present in the sample?

ANSWER: 123.4 cm–1 μM–1 61. Which of the following statements best describes the principle of mass spectrometry? a. It measures the absorption of infrared radiation by a molecule. b. It analyzes the fragmentation pattern of a molecule. c. It determines the chemical shift of protons in a molecule. d. It measures the mass-to-charge ratio of ions. e. It measures the molar mass of molecules.

ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 22 Chemical Analysis 62. In a mass spectrum produced using electron ionization, the peak with the highest m/z value is usually the a. fragment ion. b. molecular ion. c. base peak. d. radical ion. e. atomic ion.

ANSWER: b 63. Smaller peaks observed in a mass spectrum are usually a. fragment ions. b. molecular ions. c. base peaks. d. radical ions. e. atomic ions.

ANSWER: a 64. Which of the following equations represents the mechanism to produce an ion using electron ionization? a. M + OH– → MOH– b. M + NH4+ → M+ + NH3 c. M+ + e– → M d. M + e– → M+ + 2e– e. M + hν → M+ ANSWER: d 65. Which of the following equations represents the mechanism to produce an ion using chemical ionization? a. M + OH– → MOH– b. M + NH4+ → M+ + NH3 c. M+ + e– → M d. M + e– → M+ + 2e– e. M + hν → M+ ANSWER: b 66. What is one reason to use electron ionization instead of chemical ionization when performing chemical analysis using a mass spectrometer? a. Electron ionization is a softer ionization technique than chemical ionization. b. Electron ionization leads to less fragmentation than chemical ionization.

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Chapter 22 Chemical Analysis c. Electron ionization produces a higher abundance of molecular ions. d. Electron ionization produces more fragment ions, which can be useful for structural determination. e. Electron ionization preferentially ionizes polar molecules.

ANSWER: d 67. What is the "fingerprint region" in an infrared (IR) spectrum? a. The region between 3500–2800 cm–1 where O–H and C–H stretches occur. b. The region between 2000–1500 cm–1 where C=O and C=C stretches occur. c. The region between 1200–400 cm–1 which contains multiple peaks. d. The region where the most intense peak in the spectrum is located. e. The region with the widest peak.

ANSWER: c 68. Which of the following functional groups has a characteristic broad absorption peak around 3300 cm–1 in an IR spectrum? a. carbonyl group (C=O) b. hydroxyl group (O– H) c. alkene group (C=C) d. amine group (O–H) e. ether group (C–O–C) ANSWER: b 69. What is the purpose of the external magnetic field in NMR spectroscopy? a. To ionize the sample. b. To fragment the molecule. c. To generate radio waves. d. To cause the nuclear spins to align. e. To cause the sample to fluoresce.

ANSWER: d 70. In an NMR spectrum a proton with three neighboring protons will have its signal split into how many peaks? a. 2 b. 3 c. 4 d. 5 e. 6 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 22 Chemical Analysis ANSWER: c 71. What is the integration value in an 1H NMR spectrum proportional to?

ANSWER: The integration value is proportional to the number of protons giving rise to the signal.

72. In an NMR spectrum, a hydrogens signal is split into five peaks. What information does that provide?

ANSWER: The splitting pattern indicates that the hydrogen has four neighboring hydrogen atoms.

73. How many peaks would be observed in an NMR spectrum for the signal produced by the H marked by the asterisk?

ANSWER: That hydrogen would produce a signal with five peaks.

74. Describe the NMR spectrum, in terms of integration and splitting, that would be produced by acetic acid, CH3COOH. ANSWER The NMR spectrum would contain two peaks. One peak would be three times : larger, produced by hydrogens bonded to carbon in the CH3 group. The smaller peak would be produced by the hydrogen bonded to oxygen in the carboxylic acid group.

75. Describe the NMR spectrum, in terms of integration and splitting, that would be produced by ethanol, CH3CH2OH. ANSWER The NMR spectrum would contain three peaks. The signal produced by the : hydrogens in the CH3 group would have an integration of three and be split into three peaks. The signal produced by the hydrogens in the CH2 group would have an integration of two and be split into five peaks. The signal produced by the hydrogen in the OH group would have an integration of one and be split into three peaks.

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Chapter 21 Organic Chemistry and Biomolecules 1. What is studied in the field of organic chemistry? a. Earth's carbon cycle b. chemical reactions of living things c. carbon-containing molecules d. molecules associated with cells e. photosynthesis and cellular respiration ANSWER: c 2. Which of the following is a key reaction of the carbon cycle? a. condensation b. photosynthesis c. polymerization d. cellular respiration e. both b and d ANSWER: e 3. Which of the following is NOT a process involved in the carbon cycle? a. photosynthesis b. cellular respiration c. decomposition d. polymerization e. combustion ANSWER: d 4. What is the carbon cycle?

ANSWER The carbon cycle describes the different processes in which Earth’s carbon : moves between rock and sediment, water and atmosphere, and plants and animals.

5. How many covalent bonds does carbon form in a stable neutral molecule? a. 1 b. 2 c. 3 d. 4 e. It depends on which elements it is bonding with.

ANSWER: d 6. What is the condensed formula for 1-pentanol?

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Chapter 21 Organic Chemistry and Biomolecules

a. CH3(CH2)3OH b. CH3(CH2)4OH c. CH3(CH2)5OH d. CH3(CH3)3OH e. CH3(CH3)4OH ANSWER: b 7. What is the molecular formula for the following compound?

a. C6H14 b. C7H16 c. C8H18 d. C7H14 e. C8H16 ANSWER: c 8. How many constitutional isomers have a molecular formula of C6H14? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: e 9. What is the molecular formula for the amino acid alanine?

a. C3H8NO2 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 21 Organic Chemistry and Biomolecules b. C3H7NO2 c. C3H9NO2 d. C4H7NO2 e. C34H8NO2 ANSWER: b 10. What is the molecular formula for benzene?

a. C6H14 b. C6H12 c. C6H10 d. C6H8 e. C6H6 ANSWER: e 11. Which of the following statements describes the relationship between these two molecules?

a. They are the same molecule. b. They are geometric isomers. c. They are constitutional isomers. d. They are enantiomer pairs. e. They are diastereomers.

ANSWER: c 12. What are geometric isomers that are not mirror images of each other called? a. They are the same molecule. b. chiral molecules c. constitutional isomers d. enantiomer pairs e. diastereomers ANSWER: e Copyright Macmillan Learning. Powered by Cognero.

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Chapter 21 Organic Chemistry and Biomolecules 13. Is the following molecule in the cis or trans configuration?

ANSWER: The molecule is in the cis configuration.

14. Is the following molecule in the cis or trans configuration?

ANSWER: The molecule is in the trans configuration.

15. What are enantiomers?

ANSWER Enantiomers are a pair of molecules that are non-superimposable mirror : images of each other.

16. Draw the Lewis structures for the two constitutional isomers that have the molecular formula C2H6O. ANSWER:

17. Draw and label the structures of the cis and trans isomers with the molecular formula C2H2I2, which contain a double bond between the carbon atoms.

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Chapter 21 Organic Chemistry and Biomolecules ANSWER:

trans

cis

18. What class of hydrocarbons contain only single bonds? a. alkynes b. alkenes c. alkanes d. unsaturated hydrocarbons e. aromatics ANSWER: c 19. What is the name of the hydrocarbon with the condensed formula CH3(CH2)4CH3? a. heptane b. hexane c. pentane d. butane e. cyclopentane ANSWER: b 20. What is the molecular formula for the alkane octane? a. C8H16 b. C8H15 c. C8H14 d. C8H17 e. C8H18 ANSWER: e 21. What is the molecular formula for cyclopentane? a. C5H12 b. C5H11 c. C5H10 d. C5H8 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 21 Organic Chemistry and Biomolecules e. C5H5 ANSWER: c 22. What class of hydrocarbon does the following molecule belong to?

a. alkyne b. alkene c. alkane d. cycloalkane e. aromatic ANSWER: a 23. What functional group does CH3OH contain? a. aldehyde b. carbonyl c. ketone d. alcohol e. carboxylic acid ANSWER: d 24. What functional group does the following molecule contain?

a. aldehyde b. carbonyl c. ketone d. alcohol e. carboxylic acid ANSWER: e 25. Which functional group contains a carbonyl bonded to a carbon atom on one side and a hydrogen atom on Copyright Macmillan Learning. Powered by Cognero.

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Chapter 21 Organic Chemistry and Biomolecules the other? a. aldehyde b. carbonyl c. ketone d. alcohol e. carboxylic acid ANSWER: a 26. Which functional group contains a carbonyl bonded to carbon atoms on both sides? a. aldehyde b. carbonyl c. ketone d. alcohol e. carboxylic acid ANSWER: c 27. What functional group does the following molecule contain?

a. aldehyde b. ester c. ketone d. ether e. carboxylic acid ANSWER: b 28. What functional groups does the following molecule contain?

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Chapter 21 Organic Chemistry and Biomolecules a. carboxylic acid and cycloalkane b. ester and aromatic group c. ketone, alcohol, and aromatic group d. carbonyl, alcohol, and aromatic group e. carboxylic acid and aromatic group ANSWER: e 29. What functional group contains a carbonyl attached to a nitrogen atom? a. aldehyde b. amine c. ketone d. amide e. carboxylic acid ANSWER: d 30. What functional group does the molecule CH3CH2NH2 contain? a. amide b. amine c. ketone d. alcohol e. carboxylic acid ANSWER: b 31. What functional group does the following molecule contain?

a. aldehyde b. amine c. amide d. ester e. carboxylic acid ANSWER: c 32. Draw the corresponding Lewis structure for the following molecule.

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Chapter 21 Organic Chemistry and Biomolecules

ANSWER:

33. What functional groups are present in the amino acid tyrosine?

ANSWER Tyrosine contains an alcohol, an aromatic ring, an amine, and a carboxylic acid : functional group.

34. What functional groups are present in pyruvic acid?

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Chapter 21 Organic Chemistry and Biomolecules

ANSWER: Pyruvic acid contains a ketone and a carboxylic acid functional group.

35. What are the key differences between alkanes, alkenes, and alkynes?

ANSWER Alkanes, alkenes, and alkynes are all hydrocarbons containing only carbon and : hydrogen atoms. Alkanes consist of carbon atoms connected by single bonds.

Alkenes have at least one carbon-carbon double bond. Alkynes contain at least one carbon–carbon triple bond.

36. What is the structural difference between ketones and aldehydes?

ANSWER Ketones and aldehydes both contain a carbonyl functional group. For ketones, : the carbonyl is bonded to two carbon atoms. For aldehydes, the carbonyl is bonded to one carbon and one hydrogen atom.

37. Draw the skeletal structure of cyclohexane.

ANSWER:

38. What functional groups are present in the asthma medicine albuterol?

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Chapter 21 Organic Chemistry and Biomolecules

ANSWER Albuterol contains three alcohol groups, an aromatic group, and an amine : functional group.

39. What functional group(s) are present in urea?

ANSWER: Urea contains the amide functional group.

40. Draw the skeletal structure of benzene, an aromatic compound with the molecular formula C6H6.

ANSWER:

41. What type of reaction occurs when atoms or groups of atoms are added to a molecule containing a pi bond, replacing the pi bond with two new substituents? a. addition reaction b. elimination reaction c. condensation reaction d. hydrolysis reaction e. substitution reaction Copyright Macmillan Learning. Powered by Cognero.

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Chapter 21 Organic Chemistry and Biomolecules ANSWER: a 42. What type of reaction occurs when two smaller molecules combine to produce water plus a larger molecule? a. addition reaction b. elimination reaction c. condensation reaction d. hydrolysis reaction e. substitution reaction ANSWER: c 43. What type of molecule is produced when a condensation reaction occurs between a carboxylic acid and an alcohol? a. ether b. ester c. amide d. aldehyd e e. ketone ANSWER: b 44. What type of molecule is produced when a condensation reaction occurs between two alcohol molecules? a. aldehyd e b. ester c. amide d. ether e. ketone ANSWER: d 45. What type of reaction is represented by the following chemical equation?

C2H4 + I2 → C2H4I2 a. substitution reaction b. elimination reaction c. condensation reaction d. hydrolysis reaction e. addition reaction ANSWER: e 46. What type of reaction is represented by the following chemical equation?

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Chapter 21 Organic Chemistry and Biomolecules

a. substitution reaction b. elimination reaction c. condensation reaction d. hydrolysis reaction e. addition reaction ANSWER: b 47. What type of reaction is described by the following statement?

In the presence of concentrated acids, an alcohol produces a water molecule and an alkene. a. substitution reaction b. elimination reaction c. condensation reaction d. hydrolysis reaction e. addition reaction ANSWER: b 48. What is always a product of condensation reactions? a. an alcohol b. a substituted alkane c. a water molecule d. an acid e. a strong base ANSWER: c 49. Which two molecules are needed to produce the following compound during a condensation reaction?

a. two methanol molecules b. two ethanol molecules c. ethanol and acetic acid d. methanol and acetic acid Copyright Macmillan Learning. Powered by Cognero.

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Chapter 21 Organic Chemistry and Biomolecules e. ethanol and methanol ANSWER: e 50. What type of molecule is produced when a condensation reaction occurs between a carboxylic acid and an amine? a. aldehyd e b. ester c. amide d. ether e. ketone ANSWER: c 51. Predict the products of the following substitution reaction.

CH3—Cl + HS– → ANSWER: CH3—SH + Cl– 52. Predict the products of the following reaction.

ANSWER:

53. What molecule would be produced if the following molecule reacted with hydrogen in the presence of a palladium catalyst?

ANSWER:

54. Predict the products of the following condensation reaction.

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Chapter 21 Organic Chemistry and Biomolecules ANSWER:

55. What are molecules composed of long chains of simple repeating units connected by covalent bonds called? a. aromatics b. monomers c. carbohydrates d. polymers e. hydrocarbons ANSWER: d 56. What type of reaction is used to produce polymers like polyester? a. addition b. elimination c. condensation d. hydrolysis e. substitution ANSWER: c 57. What is the term for the small molecule used to produce the long chains in plastics? a. dimer b. monosaccharid e c. polymer d. polyvinyl e. monomer ANSWER: e 58. What type of plastic is made from ethylene?

ANSWER: polyethylene 59. What monomer is used to make PVC, poly(vinyl chloride)?

ANSWER: vinyl chloride 60. What type of molecule is used to store energy in plants? a. proteins b. carbohydrates c. lipids Copyright Macmillan Learning. Powered by Cognero.

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Chapter 21 Organic Chemistry and Biomolecules d. nucleic acids e. amino acids ANSWER: b 61. What type of molecules are used to make proteins? a. polysaccharides b. carbohydrates c. lipids d. nucleic acids e. amino acids ANSWER: e 62. What type of molecules are used to make polysaccharides? a. proteins b. carbohydrates c. lipids d. nucleic acids e. amino acids ANSWER: b 63. What type of molecules are primarily nonpolar and do not dissolve in water? a. proteins b. carbohydrates c. lipids d. nucleic acids e. amino acids ANSWER: c 64. What type of biomolecule is a major component of cell membranes? a. proteins b. carbohydrates c. polysaccharides d. phospholipids e. amino acids ANSWER: d 65. What biomolecule is responsible for transporting genetic information to other parts of the cell? a. proteins b. hemoglobin c. RNA Copyright Macmillan Learning. Powered by Cognero.

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Chapter 21 Organic Chemistry and Biomolecules d. DNA e. enzymes ANSWER: c 66. Which of the following molecules is NOT a base in DNA? a. thymine b. uracil c. cytosine d. adenine e. guanine ANSWER: b 67. What type of bond links amino acids together in a protein chain? a. hydrogen bond b. ionic bond c. disulfide bond d. peptide bond e. glycosidic bond ANSWER: d 68. Which of the following is NOT a component of a nucleotide? a. nitrogen containing base b. pentose sugar c. phosphate group d. amino acid e. b and c ANSWER: d 69. What type of bond holds the two strands of DNA together? a. covalent bonds b. ionic bonds c. amide linkages d. dispersion forces e. hydrogen bonds ANSWER: e 70. Which nitrogenous base pairs with adenine in DNA? a. guanine b. cytosine Copyright Macmillan Learning. Powered by Cognero.

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Chapter 21 Organic Chemistry and Biomolecules c. uracil d. thymine e. alanine ANSWER: d 71. What is the role of mRNA in protein synthesis? a. Carries amino acids to the ribosome. b. Catalyzes the formation of peptide bonds. c. Provides the genetic code for the protein. d. Forms the structural framework of the ribosome. e. Transports the protein to its final destination.

ANSWER: c 72. What are the four bases of DNA? How do they pair in DNA's double-helix structure?

ANSWER The four bases of DNA are adenine (A), thymine (T), cytosine (C), and : guanine (G). Adenine pairs with thymine, and cytosine pairs with guanine in the double helix.

73. What are the three differences in the structures of DNA and RNA? ANSWER DNA contains the sugar deoxyribose, while RNA contains the sugar : ribose. DNA has the base thymine (T), while RNA has uracil (U) instead of thymine. RNA does not adopt the double-helix arrangement of DNA. 74. What is the difference in structure between saturated and unsaturated fatty acids?

ANSWER Saturated fatty acids have only single bonds between carbon atoms in their : hydrocarbon chain, meaning each carbon is "saturated" with hydrogen atoms. Unsaturated fatty acids have one or more double bonds between carbon atoms in their hydrocarbon chain.

75. Draw the dipeptide Cys-Ser. The structures of cysteine, Cys, (left), and serine, Ser, (right) are shown.

ANSWER:

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Chapter 20 Transition Elements and Modern Materials 1. What are naturally occurring ionic compounds that are the make-up of rocks? a. metal ores b. minerals c. alloys d. pure metals e. sediment ANSWER: b 2. What are rocks or minerals that contain a specific metal? a. metal ores b. minerals c. alloys d. pure metals e. sediment ANSWER: a 3. What is the process called for using fire to extract metal from ores? a. roasting b. refining c. calcination d. smelting e. alloying ANSWER: d 4. What two metals are used to make bronze? a. copper and lead b. copper and zinc c. copper and tin d. lead and tin e. copper and iron ANSWER: c 5. Which element is combined with iron to make steel? a. copper b. zinc c. tin d. oxygen e. carbon Copyright Macmillan Learning. Powered by Cognero.

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Chapter 20 Transition Elements and Modern Materials ANSWER: e 6. What type of reaction is used to manufacture aluminum? a. decomposition b. acid dissolution c. electrolysis d. smelting e. precipitation ANSWER: c 7. When heated, solid silver carbonate decomposes into solid silver, carbon dioxide gas, and oxygen gas. Write a balanced equation for this reaction, including phase symbols. In this reaction, is silver oxidized, reduced, or neither? a. AgCO3(s) → Ag(s) + CO2(g) + O(g); silver is reduced b. AgCO3(s) → Ag(s) + CO2(g) + O(g); silver is oxidized c. Ag2CO3(s) → 2Ag(s) + CO2(g) + ½O2(g); silver is reduced d. Ag2CO3(s) → 2Ag(s) + CO2(g) + ½O2(g); silver is oxidized e. AgCO3(s) → Ag(s) + CO2(g) + ½O2(g); silver is reduced ANSWER: c 8. When heated, solid mercury(II) oxide decomposes into liquid mercury and oxygen gas. Write a balanced equation for this reaction, including phase symbols. In this reaction, is mercury oxidized, reduced, or neither? a. 2HgO(s) → 2Hg(l) + O2(g); mercury is reduced b. 2HgO(s) → 2Hg(l) + O2(g); mercury is oxidized c. 2Hg2O(s) → 4Hg(l) + O2(g); mercury is reduced d. 2Hg2O(s) → 4Hg(l) + O2(g); mercury is oxidized e. HgO2(s) → Hg(l) + O2(g); mercury is reduced ANSWER: a 9. Solid potassium chlorate decomposes upon heating to produce solid potassium chloride and oxygen gas.

Write a balanced equation for this reaction, including phase symbols. In this reaction, is chlorine oxidized, reduced, or neither?

ANSWER: 2KClO3(s) → 2KCl(s) + 3O2(g); chlorine is reduced.

10. The smelting of galena (PbS), the most common lead ore, involves several reactions. In one method, galena is roasted in the presence of oxygen. This converts the lead(II) sulfide into lead(II) oxide while producing sulfur dioxide gas. In the next step, the lead(II)oxide is reduced by heating it in the presence of solid carbon. This reduces the lead oxide to molten lead metal and produces carbon dioxide gas. Write out the balanced chemical equations for this process.

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Chapter 20 Transition Elements and Modern Materials ANSWER: (1) 2PbS(s) + 3O2(g) → 2PbO(s) + 2SO2(g) (2) 2PbO(s) + C(s) → 2Pb(l) + CO2(g) 11. Why did humans start producing iron and steel later than copper and bronze?

ANSWER Iron has a higher melting point than copper, which requires more advanced : metallurgy.

12. Molybdenum can be obtained from the mineral molybdenite, MoS2. What is the formal chemical name for this mineral? What is the oxidation number for molybdenum in molybdenite?

ANSWER: molybdenum(IV) sulfide, +4 13. Which of the following metals has the highest metallic character? a. gold b. silver c. lead d. copper e. calcium ANSWER: e 14. Rank the following elements in order of increasing metallic character: Cr, Si, Zn, Al. a. Cr < Zn < Al < Si b. Si < Al < Zn < Cr c. Al< Zn < Si < Cr d. Cr < Zn < Si < Al e. Si < Zn < Al < Cr ANSWER: b 15. Which period 4 metal has the smallest atomic radius? a. Ca b. Sc c. Co d. V e. K ANSWER: c 16. How many outer electrons does a neutral tungsten atom have? a. 2 b. 6 c. 14 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 20 Transition Elements and Modern Materials d. 18 e. 20 ANSWER: e 17. What is the electron configuration for Tc? a. [Kr]5s24d5 b. [Kr]5s14d6 c. [Kr]4d7 d. [Kr]5s24f5 e. [Kr]5s14f6 ANSWER: a 18. Molybdenum preferentially half-fills its orbitals. What is the electron configuration for Mo? a. [Kr]5s24d4 b. [Kr]5s14d5 c. [Kr]4d6 d. [Kr]5s24f4 e. [Kr]5s14f5 ANSWER: b 19. What is the term for a substance with all its electrons paired together? a. ferromagnetic b. paramagnetic c. diamagnetic d. permanent magnet e. magnetic dipole ANSWER: c 20. How many outer electrons does a Pb2+ cation have? a. 2 b. 30 c. 10 d. 26 e. 28 ANSWER: d 21. What is the electron configuration of zinc? Is zinc paramagnetic or diamagnetic?

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Chapter 20 Transition Elements and Modern Materials ANSWER: [Ar]4s23d10; diamagnetic 22. What is the electron configuration of lead? Is lead paramagnetic or diamagnetic?

ANSWER: [Xe]6s24f145d106p2; paramagnetic 23. What is metallic character? What is the periodic trend for metallic character?

ANSWER Metallic character is a measure of an element's ability to lose or share : electrons. Metallic character increases from right to left in a period and top to bottom in a group.

24. What property of metal lattices enables metals to conduct electricity efficiently?

ANSWER Metal lattices do not have fixed charges, allowing electrons to move easily : between atoms.

25. What two physical properties are caused by metals and alloys having close-packed arrangements of their lattices?

ANSWER This close packing of metal atoms accounts for metals' high density and their : ability to conduct heat efficiently.

26. What is a ferromagnetic metal?

ANSWER: A ferromagnetic metal is a metal that is able to form a permanent magnet.

27. What is the molecular geometry with a coordination number of five? a. linear b. trigonal planar c. tetrahedral d. trigonal bipyramidal e. octahedral ANSWER: d 28. What is the molecular geometry of the complex ion, [Fe(H2O)6]2+? a. linear b. trigonal planar c. tetrahedral d. trigonal bipyramidal e. octahedral ANSWER: e 29. What is the coordination number for cobalt in the complex ion, [Co(NH3)5Cl]2+? a. 2 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 20 Transition Elements and Modern Materials b. 3 c. 4 d. 5 e. 6 ANSWER: e 30. What is the molecular geometry of the complex ion in the compound, K[Ag(CN)2]? a. linear b. trigonal planar c. tetrahedral d. trigonal bipyramidal e. octahedral ANSWER: a 31. What is the name of the complex ion, [Cr(H2O)6]3+? a. hexahydroxochromium(III) ion b. hexahydratechromium(III) ion c. hexaaquachromium(III) ion d. triaquachromium(VI) ion e. triahydrochromium(VI) ion ANSWER: c 32. What is the empirical formula for diamminesilver(I) chloride? a. [Ag(CN)2]Cl b. [Ag(NH3)2]Cl c. [Ag(Cl)2]NH3 d. [Ag(Cl)2]CN e. [AgNH3]Cl2 ANSWER: b 33. What is the empirical formula for potassium tetrahydroxozincate(II)? a. K[Zn(OH)4] b. K2[Zn(OH)4] c. K2[Zn(H2O)4] d. K2[Zn(OH)2] Copyright Macmillan Learning. Powered by Cognero.

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Chapter 20 Transition Elements and Modern Materials e. K[Zn(H2O)4] ANSWER: b 34. What is the name of the following complex ion?

a. pentaaminenickelate(III) b. pentacyanonickelate(II) c. pentacarbonylnickelate(II) d. pentacyanonickelate(III) e. pentacarbonylnickelate(III) ANSWER: b 35. What is the term for the formation of polydentate bonds to a metal ion? a. precipitation b. complexation c. chelation d. alloying e. dendation ANSWER: c 36. What is the term for a ligand that has two bonding sites? a. bilagand b. diligand c. bipyramidal d. bidentate e. tridentate ANSWER: d 37. What is the formula and coordination number for calcium tetrachlorocopper(II)?

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Chapter 20 Transition Elements and Modern Materials ANSWER: Ca[CuCl4]; 4 38. What is the formula for pentacarbonyliron(0)?

ANSWER: Fe[CO]5 39. What is the name and molecular geometry of the complex ion, [NbCl5]2–?

ANSWER: pentachloroniobate(III); trigonal bipyramidal 40. What is a coordination complex?

ANSWER A coordination complex is an ion or molecule that contains a central metal ion : that is bound to multiple electron-donating ligands.

41. Draw the complex ion with the following formula: [Cu(H2O)6]2+.

ANSWER:

42. Which of the following is the weakest field ligand? a. Br– b. H2O c. OH– d. CN– e. NH3 ANSWER: a 43. Which of the following is the strongest field ligand? a. Br– b. H2O c. Cl– d. CN– e. F– ANSWER: d 44. What statement describes the arrangement of d orbitals in an octahedral environment?

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Chapter 20 Transition Elements and Modern Materials a. All five d orbitals have the same energy. b. One d orbital has higher energy. c. Two d orbitals have higher energy. d. Three d orbitals have higher energy. e. Four d orbitals have higher energy.

ANSWER: c 45. A low-spin complex is likely to form when a. the metal ion has a small charge. b. the ligands are weak-field ligands. c. the ligands are strong-field ligands. d. the complex is tetrahedral. e. the complex is not colored.

ANSWER: c 46. A high-spin complex is likely to form when a. the metal ion has a large charge. b. the ligands are weak-field ligands. c. the ligands are strong-field ligands. d. the complex is square planar. e. the complex absorbs light at a short wavelength.

ANSWER: b 47. The complex [Co(NH3)6]3+ has a maximum absorbance at 470 nm. What is the difference in energy between the lower- and higher-energy d orbitals? a. 5.1×10–19 J b. 7.0×10–19 J c. 2.3×10–19 J d. 4.2×10–19 J e. 1.9×10–19 J ANSWER: d 48. The complex [Cu(H2O)6]2+ has a d-d energy gap of 3.21×10–19 J. What would be the corresponding maximum absorbance associated with the energy difference? a. 536 nm b. 705 nm c. 425 nm d. 557 nm Copyright Macmillan Learning. Powered by Cognero.

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Chapter 20 Transition Elements and Modern Materials e. 619 nm ANSWER: e 49. How many unpaired electrons would a d6 metal ion have in a high-spin octahedral field? a. 0 b. 1 c. 2 d. 4 e. 5 ANSWER: d 50. How many unpaired electrons would a d6 metal ion have in a low-spin octahedral field? a. 0 b. 1 c. 2 d. 4 e. 5 ANSWER: a 51. The copper(II) complex ion has a maximum absorbance at 525 nm. What is the difference in energy between the lower- and higher-energy d orbitals?

ANSWER: 3.79×10–19 J 52. What is the λmax (in nm) for a complex ion in an octahedral environment that has a d-d energy gap of 2.99×10–19 J? ANSWER: 665 nm 53. What causes the splitting of d orbitals in an octahedral field?

ANSWER In an octahedral environment, the two orbitals that lie directly on the x, y, and z : axes are closer to the electron pairs of the ligands. This raises the energy of the two orbitals. The remaining three orbitals, which are farther from the ligands, are lower in energy.

54. A low-spin complex contains the Mn2+ ion. Sketch the d-electron configuration for this ion and identify whether it is paramagnetic or diamagnetic.

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Chapter 20 Transition Elements and Modern Materials ANSWER:

The complex is paramagnetic.

55. What type of elements exhibit properties between those of metals and nonmetals? a. conductors b. transition metals c. metalloids d. insulators e. complex ions ANSWER: c 56. Which of the following elements can be used as a semiconductor? a. gold b. chromiu m c. lithium d. sulfur e. silicon ANSWER: e 57. Which of the following elements would you expect to have a small bandgap? a. silver b. chromium c. sodium d. germanium e. carbon ANSWER: d 58. Which of the following elements would you expect to have a large bandgap? a. gold b. chromium c. aluminum d. germanium e. carbon Copyright Macmillan Learning. Powered by Cognero.

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Chapter 20 Transition Elements and Modern Materials ANSWER: e 59. What band contains the highest-energy orbitals that are filled? a. valence band b. conduction band c. band gap d. insulation band e. semiconductor band ANSWER: a 60. What band consists of empty orbitals in the ground state? a. valence band b. conduction band c. band gap d. insulation band e. semiconductor band ANSWER: b 61. What process is used to create patterns on a semiconductor using UV light? a. laser ablation b. photospectroscopy c. photovoltaic cells d. photolithography e. phototransistor ANSWER: d 62. What are the differences in bandgap between conductors, insulators, and semiconductors? How does this impact the conductance of each type of substance?

ANSWER Conductors have no bandgap because their valence and conduction bands : overlap. This allows electrons to easily move from the valence band to the conduction band, which is why conductors readily carry electrical current. Insulators have a large bandgap. The large energy gap makes it difficult for electrons to jump from the valence band to the conduction band, effectively blocking the flow of electricity and making them poor conductors. Semiconductors have a small bandgap, which allows some electrons to jump to the conduction band. This makes them moderately conductive.

63. What are the four general steps of the photolithography process?

ANSWER: 1. Coat the surface with photoresist.

2. Shine a light pattern on the photoresist coating.

3. Use an etching agent that removes surfaces where there is no Copyright Macmillan Learning. Powered by Cognero.

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Chapter 20 Transition Elements and Modern Materials photoresist.

4. Remove the rest of the photoresist, leaving the patterned surface behind.

64. How are semiconductors used in computers?

ANSWER Semiconductors are used to produce transistors. Transistors act as on-off : switches and are the foundation for modern computers.

65. What is the broad category of ionic compounds or mixtures that form hard, durable surfaces when heated? a. glass b. ceramics c. concrete d. polymers e. plastics ANSWER: b 66. What type of substance is produced from an amorphous, network covalent solid of silicon dioxide? a. glass b. ceramics c. nanotubes d. polymers e. plastics ANSWER: a 67. What is the temperature at which glass changes from hard and brittle to soft and malleable? a. glass boiling point b. glass conversion point c. glass melting point d. glass transition temperature e. glass metamorphosis temperature ANSWER: d 68. What type of substance is made of long chains of covalently bonded atoms? a. fullerenes b. ceramics c. nanotubes d. polymers e. graphite ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 20 Transition Elements and Modern Materials 69. What is the term for small repeating molecules used to make plastics? a. monodentate b. monomers c. nanoparticle s d. fullerenes e. carbon fibers ANSWER: b 70. What type of substance is formed by graphite rearranging to form spherical structures? a. nanotubes b. carbon fibers c. fullerenes d. polymers e. graphene ANSWER: c 71. What carbon-based substance has no band gap? a. graphite b. carbon fibers c. fullerenes d. polymers e. graphene ANSWER: e 72. What carbon-based substance is made of graphene sheets that have rolled into a hollow cylinder? a. fullerenes b. carbon fibers c. carbon nanotubes d. polymers e. graphite ANSWER: c 73. What are the similarities and differences in the structures of graphite and graphene?

ANSWER Graphite and graphene are formed by carbon atoms forming two-dimensional : sheets in a honeycomb pattern of six-atom rings. Graphite contains many of these sheets stacked on one another, whereas graphene comprises only a few layers.

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Chapter 20 Transition Elements and Modern Materials 74. What is the field of nanotechnology?

ANSWER Nanotechnology involves producing and characterizing materials and devices : at the nanometer scale.

75. What makes graphene such a promising material for a lot of industries?

ANSWER Graphene has no bandgap, so it can conduct electricity. It is also transparent, : lightweight, flexible, and mechanically strong.

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Chapter 19 Nuclear Chemistry 1. What is the collective term used for protons and neutrons? a. nucleons b. nuclides c. isotopes d. ions e. atoms ANSWER: a 2. What is the atomic symbol for an isotope of calcium with 23 neutrons? a. b. c. d. e.

ANSWER: b 3. What is the atomic symbol for an isotope of krypton with 49 neutrons? a. b. c. d. e.

ANSWER: e 4. What element can have the atomic symbol

?

a. yttrium b. radium c. radon d. ytterbium e. zirconium ANSWER: a 5. Which of the following describes an isotope of chlorine-37? a. A nuclide with a mass number of 37 and an atomic number of 37 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry b. A nuclide with a mass number of 37 and an atomic number of 17 c. A nuclide with a mass number of 17 and an atomic number of 37 d. A nuclide with a mass number of 37 and an atomic number of 18 e. A nuclide with a mass number of 37 and an atomic number of 19 ANSWER: b 6. What nuclide is produced when a proton collides with oxygen-16? The process also produces two protons and a neutron. a. lithium-10 b. beryllium-10 c. boron-10 d. beryllium-11 e. boron-11 ANSWER: b 7. What nuclide is produced when a neutron collides with oxygen-16? The process also produces a proton. a. nitrogen-16 b. fluorine-15 c. nitrogen-17 d. oxygen-17 e. oxygen-15 ANSWER: a 8. Nuclear power plants bombard uranium-238 with three neutrons to start a nuclear reaction. The reaction produces three additional neutrons, barium-140, and another nuclide. What is the identity of the second nuclide produced? a. kryton-92 b. bromine-92 c. rubidium-93 d. krypton-98 e. bromine-93 ANSWER: d 9. What isotope is produced in the following reaction? a. hydrogen4 b. hydrogen5 c. helium-3 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry d. helium-4 e. helium-5 ANSWER: d 10. What species is needed to complete the following nuclear reaction? a. hydrogen-1 b. hydrogen-2 c. hydrogen-3 d. three neutrons e. helium-4 ANSWER: b 11. What is a nuclear reaction?

ANSWER A nuclear reaction is a reaction that changes the structure of an atomic nucleus, : such as the number of protons or neutrons found in the nucleus.

12. What is the atomic symbol for a technetium isotope with 54 neutrons?

ANSWER: 13. What is the atomic symbol for a bismuth isotope that has 126 neutrons?

ANSWER: 14. Write the nuclear equation for the reaction that occurs when beryllium-7 collides with a neutron and yields a nuclide and one proton.

ANSWER: 15. Write the nuclear equation for the reaction that occurs when helium-3 collides with a neutron and yields a nuclide and one proton.

ANSWER: 16. Which type of radioactive decay results in an increase in the atomic number of the decaying nucleus? a. alpha decay b. beta decay c. gamma decay d. positron emission e. electron capture ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry 17. Which of the following statements best describes gamma radiation? a. It is high-energy electromagnetic radiation. b. It is an anti-matter particle. c. It causes an increase in the nuclide's atomic number. d. It causes a decrease in the nuclide's atomic number. e. It consists of one proton and one neutron.

ANSWER: a 18. Which of the following statements best describes alpha decay? a. The nucleus ejects a particle composed of two protons and two neutrons. b. A neutron decays into a proton and an electron. c. The nucleus releases a high-energy photon. d. A proton decays into a neutron and positron. e. A proton and an electron combine to produce a neutron.

ANSWER: a 19. Which nuclide is produced from alpha decay of fermium-257? a. curium-253 b. curium-255 c. californium-251 d. californium-253 e. californium-255 ANSWER: d 20. Which nuclide is produced from positron emission of rubidium-82? a. Kr-82 b. Kr-83 c. Sr-82 d. Sr-83 e. Br-82 ANSWER: c 21. Which of the following statements best describes beta decay? a. The nucleus ejects a particle composed of two protons and two neutrons. b. A neutron decays into a proton and an electron. c. The nucleus releases a high-energy photon. d. A proton decays into a neutron and positron. e. A proton and an electron combine to produce a neutron.

ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry 22. What is the nuclear equation for the positron emission of gallium-68? a. b. c. d. e.

ANSWER: c 23. What is the nuclear equation for the beta decay of samarium-151? a. b. c. d. e.

ANSWER: e 24. Iodine-125 has a half-life of 59 days. How long will a sample of iodine-125 take to reduce to 6.25% of its original amount? a. 118 days b. 118 days c. 177 days d. 236 days e. 295 days ANSWER: d 25. Technetium-99 has a half-life of 6.0 hours. What is the mass of a 125-mg sample of technetium-99 after 48 hours? a. 0.25 mg b. 0.49 mg c. 0.98 mg d. 2.0 mg e. 4.0 mg ANSWER: b 26. Technetium-99 has a half-life of 6.0 hours. What is the mass of a 125 mg sample of technetium-99 after 48 hours?

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Chapter 19 Nuclear Chemistry a. 0.25 mg b. 0.49 mg c. 0.98 mg d. 2.0 mg e. 4.0 mg ANSWER: b 27. Phosphorus-32 is a radioactive isotope used in medical treatments for certain cancers. It has a half-life of 14.3 days. If a patient receives a dose containing 5.0 mg of phosphorus-32, how much will remain in their system after 60.0 days? a. 0.27 mg b. 0.39 mg c. 0.31 mg d. 0.79 mg e. 0.41 mg ANSWER: a 28. What is the nuclear equation for the positron emission of yttrium-86?

ANSWER: 29. What is the nuclear equation for the electron capture of 109Cd? ANSWER: 30. What is the nuclear equation for the beta decay of bismuth-214?

ANSWER: 31. What is the nuclear equation for the alpha decay of polonium-215?

ANSWER: 32. What is the nuclear equation for the alpha decay of 221Fr? ANSWER: 33. Strontium-90 is a radioactive isotope produced in nuclear fission with a half-life of 28.8 years. If a soil sample initially contains 100.0 g of strontium-90, how much will remain after 125 years?

ANSWER: 4.94 g 34. Sodium-24 is used in medicine to study blood circulation. It has a half-life of 15 hours. How long would a 750.0 mg sample of sodium-24 take to be reduced to 10.0 mg?

ANSWER: 93 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry hours 35. Which of the following instrument does not measure radioactivity? a. spectrophotometer b. dosimeter c. Geiger counter d. scintillation counter e. semiconductor counter ANSWER: a 36. Which of the following devices can a scientist wear while working in a laboratory to monitor their radiation exposure? a. spectrophotometer b. dosimeter c. Geiger counter d. scintillation counter e. semiconductor counter ANSWER: b 37. Which of the following units measures the number of decays per second? a. sievert (Sv) b. roentgen equivalent man (rem) c. curie (Ci) d. joule (J) e. becquerel (Bq) ANSWER: e 38. What type of radiation requires the most amount of material to block? a. alpha particles b. beta particles c. gamma rays d. protons e. neutrons ANSWER: c 39. What type of radiation is the most damaging to living tissue? a. alpha particles b. beta particles c. gamma rays d. protons Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry e. neutrons ANSWER: a 40. What is the unit, becquerel (Bq), equivalent to? a. one joule of energy per kilogram of body mass b.the dosage in rads that will cause the same amount of biological injury as one rad of X-rays or gamma rays c. 3.7×1010 radioactive decays per second d.the number of decays that occur each second e. one millijoule of energy per milligram of body mass ANSWER: d 41. How does a Geiger counter measure radiation?

ANSWER A Geiger counter contains a gas-filled tube. When radiation enters the tube, it : ionizes the gas particles, producing an electron. The electric current produced by this process is then measured.

42. Rank the major classes of nuclear radiation, alpha, beta, and gamma, according to their potential to damage living tissue, from least damaging to most damaging.

ANSWER: gamma < beta < alpha 43. Why are alpha particles, despite being the most damaging to living tissue, also the easiest form of radiation to block?

ANSWER Alpha particles are the easiest form of radiation to block because of their large : mass and positive charge. These particles comprise two protons and two neutrons, making them relatively large compared to other types of radiation. As a result, they interact strongly with matter, losing energy quickly over short distances.

44. Match the radiation with the type of material that can block it.

Radiation: alpha, beta, gamma Material: thick sheet of metal, thin sheet of metal, sheet of paper ANSWER: Radiation Material alpha sheet of paper beta thin sheet of metal gamma thick sheet of metal 45. The following radioactive isotopes all undergo beta decay. The isotopes are listed with their respective halflives. Which isotope poses the most significant radiation hazard if each sample contains the same concentration? a. Ni-63, 92 years b. Mn-56, 2.6 hours Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry c. K-42, 12.4 hours d. Kr-85, 10.72 years e. I-131, 8.04 days ANSWER: b 46. The following radioactive isotopes all undergo alpha decay. The isotopes are listed with their respective half-lives. Which isotope poses the least significant radiation hazard if each sample contains the same concentration? a. thorium-230, 7.7×104 years b. radium-226, 1.6×103 years c. berkelium-247, 1.4×103 years d. plutonium-242, 3.8×105 years e. uranium-234, 2.4×105 years ANSWER: d 47. A rock sample is found to contain 0.25 mg of K-40 and 0.75 mg of Ar-40. The half-life of K-40 is 1.25 billion years. How old is the rock? a. 625 million years b. 1.25 billion c. 2.5 billion years d. 3.75 billion years e. 5.0 billion years ANSWER: c 48. A wooden artifact contains only 25% of the carbon-14 present in living trees. Given that the half-life of carbon-14 is 5,730 years, approximately how old is the artifact? a. 1,432 years b. 2,865 billion c. 5,730 years d. 11,460 years e. 22,920 years ANSWER: d 49. Iodine-123 is a radioactive isotope of iodine used in nuclear medicine imaging and has a half-life of 13.2 hours. A patient was administered 12.5 ug of the isotope. How long will it take for the isotope to decay to below 10.0% of the initial amount? a. 26.4 hours b. 43.8 hours Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry c. 39.6 hours d. 78.2 hours e. 54.3 hours ANSWER: b 50. Which of the following types of radioactivity can be used for nuclear medicine imaging? a. alpha decay b. beta decay c. positron emission d. electron capture e. All of the above ANSWER: c 51. Describe how carbon-14 dating is used to determine the age of an object derived from plant matter.

ANSWER A tiny fraction of carbon atoms are carbon-14. Carbon-14 is radioactive with a : half-life of 5,730 years. As plants grow, they take in carbon dioxide from the air, including carbon-14. When the plant dies, the carbon-14 begins to decay. Based on the amount of carbon-14 left in a sample, scientists can determine the age of substances that were derived from plants.

52. A rock sample is found to contain 0.125 mg of K-40 and 0.875 mg of Ar-40. The half-life of K-40 is 1.25 billion years. How old is the rock?

ANSWER: 3.75 billion years 53. A 125-mg sample of cesium-131, an isotope with a half-life of 9.7 days, was used for radiation treatment for a cancer patient. How much of the sample is remaining after two weeks?

ANSWER: 46.0 mg 54. What is the difference between the mass of a complete nucleus and the mass of the individual particles that comprise the nucleus? a. strong nuclear force b. mass defect c. binding energy d. band of stability e. mass-energy equivalence ANSWER: b 55. What is the energy required to break a nucleus into its substituent particles? a. strong nuclear force Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry b. mass defect c. binding energy d. band of stability e. mass-energy equivalence ANSWER: c 56. What is the binding energy of a nuclide that has a mass defect of 0.037210 u? a. 3.72×10–13 J b. 1.24×10–12 J c. 7.31×10–12 J d. 5.56×10–12 J e. 9.89×10–11 J ANSWER: d 57. What is the mass defect of a nuclide that has a binding energy of 3.74×10–12 J? a. 0.019681 u b. 0.045726 u c. 0.029120 u d. 0.036984 u e. 0.025026 u ANSWER: e 58. What is the binding energy per nucleon for C-14? Carbon-14 has a mass defect equal to 1.86288×10–28 kg. a. 1.68×10–12 J b. 1.20×10–12 J c. 1.79×10–11 J d. 4.19×10–11 J e. 9.23×10–11 J ANSWER: b 59. Use the following information to calculate the mass defect of a neutral tungsten-180 atom with a mass of 179.9467 u.

Particle Proton Neutron Electron

Mass (u) 1.007276 1.008665 0.000549

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Chapter 19 Nuclear Chemistry a. 1.6984×10–27 kg b. 3.0078×10–27 kg c. 2.5751×10–27 kg d. 3.6990×10–27 kg e. 5.2217×10–27 kg ANSWER: c 60. Use the following information to calculate the binding energy of a neutral fluorine-19 atom with a mass of 18.9984 u.

Particle Mass (u) Proton 1.007276 Neutron 1.008665 Electron 0.000549 –11 a. 2.37×10 J b. 3.12×10–11 J c. 1.89×10–11 J d. 2.08×10–11 J e. 4.97×10–11 J ANSWER: a 61. Use the following information to calculate the change in energy for the following reaction. Is energy released or absorbed?

Particle Mass (u) U-233 233.039628 Th-229 229.031755 Alpha 4.00151 –13 a. 7.57×10 J is released b. 1.29×10–13 J is released c. 1.29×10–13 J is absorbed d. 9.51×10–13 J is released e. 9.51×10–13 J is absorbed ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry 62. What is the relationship between binding energy and nuclear stability?

ANSWER Nuclear stability is proportional to the binding energy. The larger the binding : energy of a nuclide, the more stable its nucleus.

63. What is the most common decay mode for large nuclides?

ANSWER: Larger nuclides often undergo α decay.

64. Use the following information to calculate the mass defect and binding energy of a neutral K-40 atom with a mass of 39.963998 u.

Particle Proton Neutron Electron

Mass (u) 1.007276 1.008665 0.000549

ANSWER: mass defect = 5.9213×10–28 kg binding energy = 5.33×10–11 J 65. Which of the following statements correctly distinguishes nuclear fission from nuclear fusion? a. Fission requires energy input, while fusion releases energy. b.Fission releases energy, while fusion requires energy input. c. Fission involves the splitting of a heavy nucleus, while fusion involves the combining of two light nuclei. d.Fission involves the conversion of mass to energy, while fusion involves the conversion of energy to mass. e. Fission occurs in stars, while fusion is typically used in nuclear power plants.

ANSWER: c 66. Which of the following represents a nuclear fission reaction? a. b. c. d. e.

ANSWER: d 67. Which of the following represents a nuclear fusion reaction? a.

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Chapter 19 Nuclear Chemistry b. c. d. e.

ANSWER: c 68. Which of the following isotopes is commonly used as fuel in nuclear fission reactors? a. H-1 b. He-4 c. U-235 d. U-238 e. Pu-239 ANSWER: c 69. Which of the following statements best describes nuclear fission? a. The process of combining two light nuclei to form a heavier nucleus. b. The process of splitting a heavy nucleus into two or more lighter nuclei. c. The process of converting protons into neutrons within a nucleus. d. The process of emitting alpha particles from an unstable nucleus. e. The process of absorbing electrons by a nucleus to form a new element.

ANSWER: b 70. Which subatomic particle is typically used to initiate a nuclear fission chain reaction? a. proton b. electron c. neutron d. photon e. alpha particle ANSWER: c 71. What is one major environmental concern associated with nuclear fission power generation? a. depletion of fossil fuels b. greenhouse gas emissions c. thermal pollution of water bodies d. long-term storage of radioactive waste e. acid rain formation ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 19 Nuclear Chemistry 72. Use the following information to calculate the amount of energy produced, in kJ/mol of uranium, from the following reaction.

Particle Mass (u) U-235 235.04392 neutron 1.00866 Ba-140 139.910581 Kr-93 92.931130 10 a. 2.36×10 kJ/mol b. 1.98×1010 kJ/mol c. 5.50×1010 kJ/mol d. 6.07×1010 kJ/mol e. 1.66×1010 kJ/mol ANSWER: d 73. What is a nuclear fusion reaction?

ANSWER: A nuclear change in which two smaller nuclei combine to form a larger nuclei.

74. Using the table of masses, calculate the mass defect for the following nuclear reaction. How much energy does this reaction produce per mole of lithium-6 that reacts?

Particle Mass (u) neutron 1.008665 Li-6 6.015122 H-3 3.016049 He-4 4.002603 ANSWER: mass defect = 8.527×10–30 kg energy = 4.62×108 kJ/mol 75. Why is fusion energy considered a cleaner alternative to fission energy?

ANSWER: Fusion does not produce long-lived radioactive waste.

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Chapter 18 Electrochemistry 1. What is the oxidation number of nitrogen in the nitrate ion, NO3–? a. + 1 b. + 2 c. + 3 d. + 4 e. + 5 ANSWER: e 2. What is the oxidation number for carbon in the oxalate ion, C2O42–? a. + 1 b. + 2 c. + 3 d. + 4 e. + 5 ANSWER: c 3. What is the oxidation number of Mn in KMnO4? a. + 2 b. + 4 c. + 5 d. + 7 e. + 9 ANSWER: d 4. What is the oxidation number of N in NH3? a. –1 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry b. –2 c. –3 d. + 1 e. + 2 ANSWER: c 5. In the reaction Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(aq), what is being reduced? a. Zn b. H+ c. Cl– d. Zn2+ e. H2 ANSWER: b 6. In the reaction 2Al(s) + 6HBr(aq) → 2AlBr3(aq) + 3H2(aq), what is being oxidized? a. Al b. H+ c. Br– d. Al3+ e. H2 ANSWER: a 7. Which species is the oxidizing agent in the following reaction?

Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g) a. Fe2O3 b. CO c. Fe d. CO2 e. This is not an oxidation-reduction reaction.

ANSWER: a 8. Which species is the reducing agent in the following reaction?

2AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + 2Ag(s) a. AgNO3 b. Cu Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry c. Cu2+ d. NO3– e. Ag ANSWER: b 9. Which of the following is NOT a redox reaction? a. 2Na(s) + Cl2(g) → 2NaCl(s) b. 2KClO3(s) → 2KCl(s) + 3O2(g) c. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) d. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) e. 2H2(g) + O2(g) → 2H2O(l) ANSWER: d 10. Which of the following is a metal displacement reaction? a. Pb(s) + 2HNO3(aq) → Pb(NO3)2(aq) + H2(g) b. Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g) c. 2C6H6(g) + 15O2(g) → 12CO2(g) + 6H2O(g) d. 2K(s) + Br2(l) → 2KBr(s) e. Mg(s) + ZnSO4(aq) → MgSO4(aq) + Zn(s) ANSWER: e 11. What is the molecular equation for the reaction between strontium metal and nitric acid?

ANSWER: Sr(s) + 2HNO3(aq) → Sr(NO3)2(aq) + H2(g) 12. What is the molecular equation for the reaction between magnesium metal and aqueous lead(II) nitrate?

ANSWER: Mg(s) + Pb(NO3)2(aq) → Mg(NO3)2(aq) + Pb(s) 13. What is the equation for the reaction between barium metal with fluorine gas?

ANSWER: Ba(s) + F2(g) → BaF2(s) 14. What are the oxidation states for each atom in K2O2?

ANSWER: K = +1 O = –1 15. What are the oxidation states for each atom in PO43–?

ANSWER: P = +5 O = –2 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry 16. After balancing the following reaction, what is the coefficient for Ag+?

Al(s) + Ag+(aq) → Ag(s) + Al3+(aq) a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: c 17. After balancing the following reaction, what is the coefficient for Fe(s)?

Ca(s) + Fe3+(aq) → Fe(s) +Ca2+(aq) a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: b 18. After balancing the following reaction, how many electrons are transferred in this reaction?

Al(s) + Fe2+(aq) → Fe(s) + Al3+(aq) a. 2 b. 3 c. 4 d. 5 e. 6 ANSWER: e 19. After balancing the following reaction in acidic conditions, describe the amount of water molecules involved in the reaction.

Cr2O72–(aq) + I–(aq) → Cr3+(aq) + I2(aq) a. 2 moles of H2O are needed as reactants. b. 5 moles of H2O are needed as reactants. c. 4 moles of H2O are produced. d. 7 moles of H2O are produced. e. 6 moles of H2O are produced.

ANSWER: d 20. After balancing the following reaction in acidic conditions, describe the amount of H+ ions involved in the Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry reaction.

NO3–(aq) + Cu(s) → NO2(g) + Cu2+(aq) a. 2 moles of H+ are needed as reactants. b. 4 moles of H+ are needed as reactants. c. 6 moles of H+ are needed as reactants. d. 7 moles of H+ are produced. e. 6 moles of H+ are produced.

ANSWER: b 21. How many electrons are transferred in the following reaction in acidic conditions?

SO2(aq) + Cr2O72−(aq) → SO42−(aq) + Cr3+(aq) a. 2 b. 3 c. 4 d. 5 e. 6 ANSWER: e 22. After balancing the following reaction in basic conditions, describe the amount of water molecules involved in the reaction.

Cr(OH)3(aq) + ClO–(aq) → CrO42–(aq) + Cl–(aq) a. 3 moles of H2O are needed as reactants. b. 4 moles of H2O are needed as reactants. c. 5 moles of H2O are produced. d. 6 moles of H2O are produced. e. 8 moles of H2O are produced.

ANSWER: a 23. After balancing the following reaction in basic conditions, describe the amount of hydroxide ions involved in the reaction.

Cl2(aq) + S2–(aq) → Cl–(aq) + SO42–(aq) a. 4 moles of OH– are needed as reactants. b. 6 moles of OH– are needed as reactants. c. 8 moles of OH– are needed as reactants. d. 3 moles of OH– are produced. e. 5 moles of OH– are produced.

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Chapter 18 Electrochemistry ANSWER: c 24. How many electrons are transferred in the following reaction that occurs in basic conditions?

Al(s) + NO3–(aq) → Al(OH)4–(aq) + NH3(aq) a. 6 b. 8 c. 24 d. 18 e. 12 ANSWER: c 25. After balancing the following reaction in acidic conditions, describe the amount of H+ ions involved in the reaction.

Cr2O72–(aq) + I–(aq) → Cr3+(aq) + I2(aq) a. 14 moles of H+ are needed as reactants. b. 10 moles of H+ are needed as reactants. c. 8 moles of H+ are needed as reactants. d. 12 moles of H+ are produced. e. 6 moles of H+ are produced.

ANSWER: a 26. Balance the following redox reaction.

Co(s) + H+(aq) → Co3+(aq) + H2(g) ANSWER: 2Co(s) + 6H+(aq) → 2Co3+(aq) + 3H2(g) 27. Balance the following redox reaction in acidic conditions.

H2O2(aq) + MnO4−(aq) → O2(aq) + Mn2+(aq) ANSWER: 5H2O2(aq) + 2MnO4−(aq) + 6H+(aq) → 5O2(aq) + 2Mn2+(aq) + 8H2O(l) 28. Balance the following redox reaction in acidic conditions.

NO2−(aq) + MnO4−(aq) → NO3−(aq) + Mn2+(aq) + H2O(l) ANSWER: 5NO2−(aq) + 2MnO4−(aq) + 6H+(aq) → 5NO3−(aq) + 2Mn2+(aq) + 3H2O(l) 29. Balance the following redox reaction in basic conditions.

SO2(aq) + MnO4−(aq) → SO42−(aq) + MnO2(s) ANSWER: 3SO2(aq) + 2MnO4−(aq) + 4OH−(aq) → 3SO42−(aq) + 2MnO2(s) + 2H2O(l) 30. Balance the following redox reaction in basic conditions.

Ag2O(s) + N2H4(aq) → Ag(s) + N2(g) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry ANSWER: Ag2O(s) + N2H4(aq) + 2OH−(aq) → 2Ag(s) + N2(g) + 3H2O(l) 31. In a voltaic cell, the cathode is a. the electrode where reduction occurs. b. the electrode where oxidation occurs. c. the site of electron production. d. the oxidizing agent. e. the source of spectator ions.

ANSWER: a 32. In a voltaic cell, the anode is a. the electrode where reduction occurs. b. the electrode where oxidation occurs. c. the site of electron production. d. the oxidizing agent. e. the source of spectator ions.

ANSWER: b 33. The component in a voltaic cell that contains spectator ions to keep each half-cell charge neutral is known as a(n) a. electrode. b. cathode. c. anode. d. salt bridge. e. voltmeter.

ANSWER: d 34. Which species is the cathode in the following cell notation?

Cu(s) | Cu2+(aq) || Br2(aq), Br–(aq) | Pt(s) a. Cu(s) b. Cu2+(aq) c. Br2(aq) d. Br–(aq) e. Pt(s) ANSWER: e 35. Which species is the anode in the following cell notation?

Cu(s) | Cu2+(aq) || Br2(aq), Br–(aq) | Pt(s) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry a. Cu(s) b. Cu2+(aq) c. Br2(aq) d. Br–(aq) e. Pt(s) ANSWER: a 36. Which species is the oxidizing agent in the following cell notation?

Cu(s) | Cu2+(aq) || Br2(aq), Br–(aq) | Pt(s) a. Cu(s) b. Cu2+(aq) c. Br2(aq) d. Br–(aq) e. Pt(s) ANSWER: c 37. What is the cell notation for the following redox reaction?

Zn(s) + Pb2+(aq) → Pb(s) + Zn2+(aq) a. Zn(s) | Pb2+(aq) || Pb(s) | Zn2+(aq) b. Pb(s) | Pb2+(aq) || Zn2+(aq) | Zn(s) c. Zn(s) | Zn2+(aq) || Pb2+(aq), Pb(s) | Pt(s) d. Zn(s) | Zn2+(aq) || Pb2+(aq) | Pb(s) e. Pb(s) | Zn2+(aq) || Pb2+(aq) | Zn(s) ANSWER: d 38. What is the cell notation for the following redox reaction?

Ca(s) + Cu2+(aq) → Cu(s) + Ca2+(aq) a. Ca(s) | Cu2+(aq) || Cu(s) | Ca2+(aq) b. Ca(s) | Ca2+(aq)|| Cu2+(aq) | Cu(s) c. Ca(s) | Ca2+(aq) || Cu2+(aq), Cu2+(aq) | Pt(s) d. Cu(s) | Cu2+(aq) || Ca2+(aq) | Ca(s) e. Cu(s) | Ca2+(aq) || Cu2+(aq) | Ca(s) ANSWER: b 39. Calculate the standard cell potential (E°cell) for the following reaction: 2Fe3+(aq) + 2I−(aq) → 2Fe2+(aq) + I2(s) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry given the following standard reduction potentials: Fe3+(aq) + e− → Fe2+(aq) E° = +0.771 V I2(s) + 2e− → 2I−(aq) E° = +0.536 V a. +0.235 V b. +1.307 V c. –0.235 V d. –1.307 V e. +1.536 V ANSWER: a 40. Calculate the standard cell potential (E°cell) for the following reaction: Co2+(aq) + Sn2+(aq) → Sn4+(aq) + Co(s) given the following standard reduction potentials: Sn4+(aq) + 2e− → Sn2+(aq) E° = +0.151 V Co2+(aq) + 2e− → Co(s) E° = –0.28 V a. +0.13 V b. +0.43 V c. –0.13 V d. –0.43 V e. +0.54 V ANSWER: d 41. What is a voltaic cell?

ANSWER A voltaic cell is a device that produces electric current from a redox reaction in : which the two half-reactions have been separated.

42. What are the two half-reactions that occur in a voltaic cell with the following cell notation?

Ni(s) | Ni2+(aq) || H+(aq) | H2(g) | Pt(s) ANSWER: Oxidation: Ni(s) → Ni2+(aq) + 2e– Reduction: 2H+(aq) + 2e– → H2(g) 43. What is the cell notation for a voltaic cell that uses the following redox reaction?

Cr(s) + Pb2+(aq) → Cr2+(aq) + Pb(s) ANSWER: Cr(s) | Cr2+(aq) || Pb2+(aq) | Pb(s) 44. A voltaic cell is constructed with a silver cathode in a 1.0 M aq. Ag+ solution and a nickel anode in a 1.0 M aq. Ni2+ solution. Calculate the standard cell potential (E°cell) for this cell given the following standard Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry reduction potentials.

Ag+(aq) + e− → Ag(s) Ni2+(aq) + 2e− → Ni(s)

E° = +0.800 V E° = –0.257 V

ANSWER: E°cell = +1.057 V 45. Calculate the standard cell potential (E°cell) for the following reaction: 3Cu2+(aq) + 2Al(s) → 3Cu(s) + 2Al3+(aq) given the following standard reduction potentials: Cu2+(aq) + 2e− → Cu(s) E° = +0.342 V 3+ − Al (aq) + 3e → Al(s) E° = –1.676 V ANSWER: E°cell = +2.018 V 46. What is the ΔG° for the following redox reaction?

Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s) E°cell = +1.925 V a. –186 kJ/mol b. –372 kJ/mol c. –254 kJ/mol d. –198 kJ/mol e. –312 kJ/mol ANSWER: b 47. What is E°cell for the following redox reaction?

Mg(s) + Ni2+(aq) → Mg2+(aq) + Ni(s) ΔG° = –408.0 kJ/mol a. –2.114 V b. 2.114 V c. –4.229 V d. 4.229 V e. 0.000 V ANSWER: b 48. Calculate the change in Gibbs free energy (ΔG°) for the following reaction: Ru2+(aq) + 2I−(aq) → Ru(s) + I2(s) given the following standard reduction potentials: Ru2+(aq) + 2e− → Ru(s) E° = +0.455 V − − I2(s) + e → 2I (aq) E° = +0.536 V a. +78.2 kJ/mol b. +234 kJ/mol Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry c. –156 kJ/mol d. –234 kJ/mol e. +156 kJ/mol ANSWER: e 49. What is K for the following redox reaction at 298 K? Fe(s) + 2H2O(l) → Fe2+(aq) + H2(g) + 2OH–(aq) E°cell = –0.381 V –13 a. 9.23×10 b. 4.28×10–11 c. 1.29×10–13 d. 7.76×1012 e. 4.47×1013 ANSWER: c 50. What is K for the following redox reaction at 298 K? Cd2+(aq) + 2Cr2+(aq) → Cd(s) + 2Cr3+(aq) Given the following standard reduction potentials: Cd2+(aq) + 2e− → Cd(s) E° = –0.403 V 3+ − 2+ Cr (aq) + e → Cr (aq) E° = –0.407 V a. 56.9 b. 1.37 c. 232 d. 4.12 e. 6.18 ANSWER: b 51. Which of the following cell potentials would produce the greatest amount of free energy for work? a. +1.203 V b. +0.956 V c. +0.234 V d. –0.963 V e. –1.143 V ANSWER: a 52. Which of the following systems would have the largest equilibrium constant? The cell potential for each system is provided. a. +0.706 V b. +0.956 V Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry c. +0.204 V d. –1.651 V e. –1.043 V ANSWER: b 53. Which of the following systems would have the smallest equilibrium constant? The cell potential for each system is provided. a. +0.498 V b. +1.256V c. –0.227 V d. –1.009 V e. –0.004 V ANSWER: d 54. What is the cell potential for the cell below if [Cu2+] = 0.200 M, [Zn2+] = 1.500 M, and T = 20.0 °C? Cu(s) ∣ Cu2+(aq) ∣∣ Zn2+(aq) ∣ Zn(s) Given the provided information: Cu2+(aq) + 2e− → Cu(s) E° = +0.342 V 2+ − Zn (aq) + 2e → Zn(s) E° = −0.762 V a. +1.13 V b. +1.10 V c. –1.10 V d. –0.917 V e. –0.420 V ANSWER: a 55. Which of the following would increase the cell potential for the following system?

2Al(s) + 3Ni2+(aq) → 2Al3+(aq) + 3Ni(s) E° = +1.416 V a. increasing the amount of Al(s) b. increasing the amount of Ni2+(aq) c. increasing the amount of Al3+(aq) d. decreasing the amount of Ni2+(aq) e. increasing the amount of Ni(s) ANSWER: b 56. A concentration cell is constructed with two hydrogen electrodes. One electrode is immersed in a solution with [H+] = 1.0 M, and the other electrode is immersed in a solution with [H+] = 0.10 M. Both compartments are maintained at a temperature of 25 °C. What is the cell potential of this concentration cell?

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Chapter 18 Electrochemistry a. +0.059 V b. +0.030 V c. –0.10 V d. –0.059 V e. –0.030 V ANSWER: a 57. Consider the following redox reaction occurring in an electrochemical cell at 25 °C. Zn(s) + Pb2+(aq) → Zn2+(aq) + Pb(s) E°cell = +0.636 V If the initial concentrations of [Pb2+] = 2.0 M and [Zn2+] = 0.20 M, what is the initial cell potential for this reaction? a. +0.666 V b. +0.333 V c. +0.636 V d. –0.596V e. –0.132 V ANSWER: a 58. Find the cell potential for a concentration cell with two iron electrodes. One electrode is immersed in a 1.20 M iron(II) chloride solution, and the other electrode is immersed in a 5.0 mM iron(II) chloride solution. The temperature is 25 °C. ANSWER: E = +0.070 V 59. A concentration cell is constructed with two copper electrodes and two copper(II) nitrate solutions. If the first cell has a Cu2+ concentration of 2.5 μM, what concentration is required for the second cell at the cathode to produce a cell potential of 0.10 V at 25 °C? ANSWER: 6.0×10–3 M 60. A concentration cell is constructed with two iron electrodes and two iron(II) acetate solutions. If the first cell has a Fe2+ concentration of 1.0 μM, what concentration is required for the second cell at the cathode to produce a cell potential of 0.15 V at 25 °C? ANSWER: 0.12 M 61. What are the signs of ΔG° and E°cell for a reaction that is spontaneous under standard conditions?

ANSWER: ΔG° < 0 and E°cell > 0 62. What are the signs of ΔG° and E°cell for a reaction that is nonspontaneous under standard conditions?

ANSWER: ΔG° > 0 and E°cell < 0 63. What is the primary advantage of lithium-ion batteries compared to lead-acid batteries?

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Chapter 18 Electrochemistry a. lower cost of production b. higher energy density c. greater environmental friendliness d. longer lifespan e. ability to use in both acidic and basic environments ANSWER: b 64. Which of the following is an example of an electrolysis reaction? a. The use of electric current to drive the reaction of hydrogen and oxygen to form water. b. The generation of electricity in a battery. c. The corrosion of iron in the presence of oxygen and water. d. The use of an electric current to split water into hydrogen and oxygen gas. e. The electroplating of a metal onto a surface.

ANSWER: d 65. Which of the following is NOTa key component of an alkaline battery? a. zinc anode b. manganese(IV) oxide cathode c. acidic electrolyte d. porous separator e. potassium hydroxide electrolyte ANSWER: c 66. What is the primary fuel used in most commercial fuel cells? a. natural gas b. ethanol c. hydrogen d. propane e. octane ANSWER: c 67. What is the process of coating a metal object with a thin layer of another metal using a redox reaction? a. electroplatin g b. lamination c. electrolysis d. oxidation e. corrosion ANSWER: a Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry 68. What is the term for the breakdown of metal structures by oxidation to metal ions? a. electroplatin g b. galvanization c. electrolysis d. oxidation e. corrosion ANSWER: e 69. Which of the following correctly describes the role of a sacrificial anode in corrosion prevention? a. It accelerates the oxidation of iron to prevent rust. b. It forms a protective layer of zinc oxide on the metal surface. c. It replaces electrons lost from the more expensive metal, preventing its oxidation. d. It is coated with nickel to prevent corrosion. e. It oxidizes water to prevent rust formation.

ANSWER: c 70. In the chlor-alkali process, what product is formed at the anode? a. Cl2 b. NaOH c. H2 d. Na+ e. O2 ANSWER: a 71. What three products are formed in the chlor-alkali process?

ANSWER: The chlor-alkali process produces chlorine gas, hydrogen gas, and sodium hydroxide.

72. What is ΔG° (in kJ/mol) for the redox reaction used in the lead-acid battery?

Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l) E°cell = 2.05 V ANSWER: ΔG° = –396 kJ/mol 73. Which type of battery is commonly used in electric vehicles (EVs) and hybrid electric vehicles (HEVs)?

ANSWER: lithium-ion batteries 74. What are the two half-reactions that occur at the cathode and anode in an alkaline battery?

Zn(s) + 2MnO2(s) → ZnO(s) + Mn2O3(s) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 18 Electrochemistry ANSWER: Anode: Zn(s) + 2OH–(aq) → ZnO(s) + H2O(l) + 2e– Cathode: 2MnO2(s) + H2O(l) + 2e– → Mn2O3(s) + 2OH–(aq) 75. What are fuel cells?

ANSWER Fuel cells are devices that convert the potential energy of a combustion : reaction directly into electrical energy by separating the oxidation and reduction half-reactions.

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Chapter 17 Thermodynamics 1. What is the first law of thermodynamics? a. In any change, the total amount of energy and entropy in the universe remains constant. b. In any change, the total amount of entropy in the universe remains constant. c. In any change, the total amount of entropy in the universe is increasing. d. In any change, the total amount of energy in the universe remains constant. e. In any change, the total amount of energy in the universe is increasing.

ANSWER: d 2. A system absorbs 10 kJ of heat from its surroundings. What is the value of qsys? a. –10 kJ b. 0 kJ c. +10 kJ d. +20 kJ e. It cannot be determined.

ANSWER: c 3. A system releases 5 kJ of heat to its surroundings. What is the value of qsurr? a. –5 kJ b. 0 kJ c. +5 kJ d. +10 kJ e. It cannot be determined.

ANSWER: c 4. Which of the following is true for any change according to the first law of thermodynamics? a. ΔEsys = ΔEsurr b. ΔEsys > ΔEsurr c. ΔEsys < ΔEsurr d. ΔEsys + ΔEsurr = 0 e. ΔEsys – ΔEsurr = 0 ANSWER: d 5. Which of the following is not a spontaneous change? a. A gas expanding to fill its container. b. Heat flowing from a hot object to a cold object. c. A ball rolling down a hill.

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Chapter 17 Thermodynamics d. Water freezing below 0 °C. e. A gas being compressed into a smaller volume.

ANSWER: e 6. What is ΔEsys if ΔEsurr = – 50 kJ? a. –50 kJ b. 0 kJ c. +50 kJ d. –100 kJ e. > 50 kJ ANSWER: c 7. What is the first law of thermodynamics?

ANSWER The first law of thermodynamics states that in any change, the total amount of : energy in the universe remains constant.

8. What is a spontaneous chemical or physical change?

ANSWER A spontaneous change is any change that happens without outside influence or : energy.

9. What is the value for ΔEsurr if ΔEsys is –250 kJ? ANSWER: ΔEsurr = +250 kJ 10. What is the relationship between ΔEsys and ΔEsur as a result of the first law of thermodynamics?

ANSWER: ΔEsys + ΔEsurr = 0 or ΔEsys = –ΔEsurr 11. Which of the following statements correctly describes the second law of thermodynamics? a. Energy is always conserved in an isolated system. b. Heat flows spontaneously from low to high temperature. c. The total entropy of the universe decreases over time. d. In any spontaneous process, the entropy of the universe increases. e. The total energy of the universe is constantly increasing.

ANSWER: d 12. According to the second law of thermodynamics, what happens to the usable energy in the universe over time? a. It remains constant. b. It increases. c. It decreases as entropy increases. d. It increases as entropy decreases.

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Chapter 17 Thermodynamics e. It depends on change in energy of the surroundings.

ANSWER: c 13. Entropy is often described as a measure of a. the energy available to do work. b. the tendency of energy to spread out or disperse. c. the heat absorbed by a system. d. the total energy within a system. e. the kinetic energy of molecules.

ANSWER: b 14. Using the Boltzmann definition, entropy can be related to a. the number of microstates in a system. b. the total energy of the system. c. the work done by the system. d. the free energy available in a system. e. the temperature of the surroundings.

ANSWER: a 15. A system contains two atoms and three units of energy. How many microstates are possible for this system? a. 2 b. 3 c. 4 d. 6 e. 8 ANSWER: c 16. A system contains four atoms and two units of energy. How many microstates are possible for this system? a. 12 b. 10 c. 8 d. 6 e. 4 ANSWER: b 17. What is the entropy of a system with one billion possible microstates? a. 1.00×10–9 J/K b. 2.86×10–22 J/K c. 4.22×10–22 J/K Copyright Macmillan Learning. Powered by Cognero.

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Chapter 17 Thermodynamics d. 1.00×10–22 J/K e. 1.13×10–22 J/K ANSWER: b 18. A system absorbs 150 mJ of heat at a constant temperature of 310 K. What is the change in entropy of the system? a. 0.48 J/K b. 2.1 J/K c. 2.7×10–3 J/K d. 4.8×10–4 J/K e. 7.2×10–3 J/K ANSWER: d 19. The surrounding air at 298 K absorbs 225 mJ of heat from a metal block at 315 K. What is the change in entropy for the metal block? a. –7.14 J/K b. +7.55×10–4 J/K c. –7.55×10–4 J/K d. +7.14×10–4 J/K e. –7.14×10–4 J/K ANSWER: d 20. What is the second law of thermodynamics?

ANSWER The second law of thermodynamics states that for any spontaneous change, the : entropy of the universe increases.

21. A system contains one mole of particles. If each particle represents a microstate, what is the entropy of the system?

ANSWER: 7.5595×10–22 J/K 22. System A has 10 microstates, and system B has 100 microstates. If both systems are at the same temperature, which system has the higher entropy?

ANSWER: Since system B has more microstates than system A, it has higher entropy.

23. What does the term free energy refer to in thermodynamics? How does the amount of free energy change during a chemical or physical change?

ANSWER Free energy is the energy available to do work in a system. During a : spontaneous chemical or physical change, the amount of free energy decreases. In a nonspontaneous change, the amount of free energy increases.

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Chapter 17 Thermodynamics 24. A system contains two atoms and four units of energy. How many microstates are possible for this system?

ANSWER: There are five possible microstates.

25. Which of the following gases would have the highest molar entropy at standard state conditions? a. H2 b. O2 c. CH4 d. SO3 e. CO2 ANSWER: d 26. Which of the following gases would have the lowest molar entropy at standard state conditions? a. F2 b. N2 c. He d. N2O4 e. NH3 ANSWER: c 27. Which of the following reactions would have a positive change in entropy (ΔS > 0)? a. N2(g) + 3H2(g) → 2NH3(g) b. CaCO3(s) → CaO(s) + CO2(g) c. 2H2(g) + 2O2(g) → 2H2O(l) d. H2O(l) → H2O(s) e. 2Zn(s) + O2(g) → 2ZnO(s) ANSWER: b 28. Which of the following reactions would have a negative change in entropy (ΔS < 0)? a. NH3(s) ⟶ NH3(l) b. NaCl(s) ⟶ Na+(aq) + Cl−(aq)

c. 2C6H14(l) + 19O2(g) ⟶14H2O(g) + 12CO2(g) d. 2Fe(s) + 3O2(g) ⟶ 2Fe2O3(s) e. H2O(l) ⟶ H2O(g)

ANSWER: d

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Chapter 17 Thermodynamics 29. According to the third law of thermodynamics, what is the entropy of a perfect crystalline solid at absolute zero? a. infinite b. equal to the number of atoms in the crystal c. dependent on the type of crystal d. equal to its enthalpy of fusion e. zero ANSWER: e 30. Which phase change results in the greatest increase in entropy? a. solid to liquid b. liquid to gas c. solid to gas d. gas to solid e. liquid to solid ANSWER: c 31. Using the following information, calculate the change in entropy for the following reaction.

H2(g) + Cl2(g) → 2HCl(g) Compound H2(g) Cl2(g) HCl(g) a. –166.9 J/K b. 20.0 J/K c. 53.7 J/K d. –76.3 J/K e. 129.3 J/K ANSWER: b

S°, J/(mol·K) 130.7 223.1 186.9

32. The melting point of aluminum is 660.3 °C. What is the entropy of fusion for aluminum if its enthalpy of fusion is 8.66 kJ/mol? a. 13.1 J/K b. 7.98 J/K Copyright Macmillan Learning. Powered by Cognero.

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Chapter 17 Thermodynamics c. 15.6 J/K d. 21.0 J/K e. 9.28 J/K ANSWER: e 33. Using the following information, calculate the change in entropy for the following reaction.

2CH3OH(l) + 3O2(g) → 2CO2(g) + 4H2O(g) Compound CH3OH(l)

S°, J/(mol·K) 126.8

O2(g)

205.2

CO2(g)

213.8

H2O(g) a. –156.9 J/K b. 209.0 J/K c. 248.7 J/K d. –252.8 J/K e. 313.6 J/K ANSWER: e

188.8

34. Using the following information, calculate the change in entropy for the following reaction.

CH3CH2OH(l) + 3O2(g) → 2CO2(g) + 3H2O(g)

CH3CH2OH(l)

S°, J/ (mol·K) 160.7

O2(g)

205.2

CO2(g)

213.8

Compound

188.8 H2O(g) ANSWER: 217.7 J/K 35. Using the following information, calculate the change in entropy for the following reaction.

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Chapter 17 Thermodynamics 2Cr(s) + 3Cl2(g) → 2CrCl3(s) Compound Cr(s) Cl2(g)

S°, J/(mol·K) 23.8 223.1

123.0 CrCl3(s) ANSWER: –470.9 J/K 36. Using the following information below, calculate the change in entropy for the following reaction.

4Al(s) + 3O2(g) → 2Al2O3(s) Compound Al(s) O2(g)

S°, J/(mol·K) 28.3 205.2

50.9 Al2O3(s) ANSWER: –627.0 J/K 37. A 25.00-g block of gallium is gradually heated at atmospheric pressure. After reaching its melting temperature of 30.0 °C, an additional 1,807 J of energy is required to melt the solid completely. What is the enthalpy of fusion (ΔHfus) in kJ/mol and the change in entropy (ΔSfus) in J/(mol·K) for this phase change?

ANSWER: ΔHfus = 5.040 kJ/mol ΔSfus = 16.90 J/(mol·K) 38. What is the third law of thermodynamics?

ANSWER The third law of thermodynamics states that at absolute zero, the entropy of a : pure crystalline solid is zero.

39. Under what conditions would a chemical reaction always be spontaneous? a. ΔH > 0 and ΔS > 0 b. ΔH < 0 and ΔS > 0 c. ΔH < 0 and ΔS < 0 d. ΔH > 0 and ΔS < 0 e. None of the above ANSWER: b 40. Under what conditions would a chemical reaction never be spontaneous? a. ΔH > 0 and ΔS > 0 b. ΔH < 0 and ΔS > 0 c. ΔH < 0 and ΔS < 0 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 17 Thermodynamics d. ΔH > 0 and ΔS < 0 e. None of the above ANSWER: d 41. Under what conditions would a chemical reaction be spontaneous at low temperatures? a. ΔH > 0 and ΔS > 0 b. ΔH < 0 and ΔS > 0 c. ΔH < 0 and ΔS < 0 d. ΔH > 0 and ΔS < 0 e. None of the above ANSWER: c 42. Calculate ΔG for the following reaction at 25 °C. Is the reaction spontaneous or nonspontaneous at that temperature?

N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = –92.38 kJ a. –124.7 kJ; spontaneous b. +4.87 kJ; spontaneous c. +4.87 kJ; nonspontaneous d. –33.3 kJ; spontaneous e. –33.3 kJ; nonspontaneous ANSWER: d

ΔS = –198.4 J/K

43. Calculate ΔG for the following reaction at 25 °C. Is the reaction spontaneous or nonspontaneous at that temperature?

2H2O2(aq) → 2H2O(l) + O2(g) ΔH = –196.1 kJ a. –199.2 kJ; nonspontaneous b. –199.2 kJ; spontaneous c. +114.9 kJ; spontaneous d. –233.4 kJ; nonspontaneous e. –233.4 kJ; spontaneous ANSWER: e

ΔS = +125.2 J/K

44. Calculate ΔG for the following reaction at 25 °C. Is the reaction spontaneous or nonspontaneous at that temperature?

CaCO3(s) → CaO(s) + CO2(g) ΔH = +178.3 kJ a. +130.5 kJ; nonspontaneous b. +130.5 kJ; spontaneous c. –227.9 kJ; spontaneous Copyright Macmillan Learning. Powered by Cognero.

ΔS = +160.3 J/K

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Chapter 17 Thermodynamics d. +174.3 kJ; nonspontaneous e. +174.3 kJ; spontaneous ANSWER: a 45. What is the boiling point of CCl4, which is the temperature at which the following transition becomes spontaneous?

CCl4(l) → CCl4(g) a. 596 K b. 728 K c. 343 K d. 272 K e. 298 K ANSWER: c

ΔH = +32.6 kJ

ΔS = +95.0 J/K

46. At what temperature does the sublimation of iodine become spontaneous?

I2(s) → I2(g) a. 497 K b. 826 K c. 643 K d. 534 K e. 298 K ANSWER: d

ΔH = +62.4 kJ

ΔS = +116 J/K

47. How much Gibbs free energy is available during the neutralization of 0.225 moles of NaOH? HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) a. 63.96 kJ b. 21.9 kJ c. 14.4 kJ d. 120 kJ e. –48.6 kJ ANSWER: c

ΔG = –63.96 kJ

48. How much Gibbs free energy is available during the combustion of 129 g of carbon?

C(s) + O2(g) → CO2(g) ΔG = –394.4 kJ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 17 Thermodynamics a. 394 kJ b. 1,210 kJ c. 987 kJ d. 4,240 kJ e. 5,080 kJ ANSWER: d 49. How much Gibbs free energy is available during the synthesis of 125 g of water at 25 °C? 2H2(g) + O2(g) → 2H2O(l) a. 879 kJ b. 1,310 kJ c. 2,710 kJ d. 3,940 kJ e. 983 kJ ANSWER: b

ΔH = –571.6 kJ

ΔS = –327 J/K

50. Determine the free energy change for the formation of ethane (C2H6) from elemental carbon and hydrogen: 2C(s) + 3H2(g) → C2H6(g) using the following reactions: C(s) + O2(g) → CO2(g) H2(g) + ½O2(g) → H2O(g) 2C2H6(g) + 3O2(g) → 2CO2(g) + 3H2O(g) ANSWER: 86.0 kJ

ΔG = –394.4 kJ ΔG = –228.6 kJ ΔG = –3121.2 kJ

51. Determine the free energy change for the formation of the ammonia, NH3, from elemental nitrogen and hydrogen: N2(g) + 3H2(g) → 2NH3(g) using the following reactions: Copyright Macmillan Learning. Powered by Cognero.

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Chapter 17 Thermodynamics N2(g) +O2(g) → 2NO(g) 2H2(g) + O2(g) → 2H2O(g) 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g) ANSWER: –166.4 kJ

ΔG = –163.2 kJ ΔG = –457.2 kJ ΔG = –1,365.2 kJ

52. At what temperature does the decomposition of calcium carbonate (CaCO3) become spontaneous?

CaCO3(s) → CaO(s) + CO2(g) ANSWER: 1,112 K

ΔH = +178.3 kJ

ΔS = +160.3 J/K

53. Calculate ΔG for the following reaction at 25 °C. Is the reaction spontaneous or nonspontaneous at that temperature?

N2O4(g) → 2NO2(g) ΔH = +58.0 kJ ΔS = +176.1 J/K ANSWER: ΔG = +5.50 kJ; Since ΔG is positive, the reaction is nonspontaneous at 25 °C. 54. Using the information below, calculate the Gibbs free energy for the following reaction.

2CH3OH(l) + 3O2(g) → 2CO2(g) + 4H2O(g) Compound

ΔGf°, kJ/mol

CH3OH(l)

–174.8

O2(g)

0

CO2(g)

–394.4

H2O(g) a. –856.9 kJ b. 1,206.2 kJ c. 1,698.7 kJ d. –1,353.6 kJ e. –913.6 kJ ANSWER: d

–228.6

55. Using the information below, calculate the Gibbs free energy or the following reaction.

P4(g) + 5O2(g) → P4O10(g) Compound

ΔGf°, kJ/mol

P4(g)

24.4

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Chapter 17 Thermodynamics O2(g)

0

P4O10(g) a. –2,697.7 kJ b. –2,722.1 k J c. –2,673.3 kJ d. –987.6 kJ e. –1,890.1 kJ ANSWER: b

–2,697.7

56. Using the information below, calculate the Gibbs free energy for the following reaction. Is the reaction spontaneous?

NH3(g) + HNO3(g) → NH4NO3(s) Compound

ΔGf°, kJ/mol

NH3(g)

–16.4

HNO3(g)

–73.5

–183.9 NH4NO3(s) a. –167.5 kJ; spontaneous b. +167.5 kJ; nonspontaneous c. –94.0 kJ; spontaneous d. +94.0 kJ; nonspontaneous e. –189.1 kJ; spontaneous ANSWER: c 57. Using the information below, calculate the Gibbs free energy for the following reaction. Is the reaction spontaneous?

4FeCl3(s) + 3O2(g) → 2Fe2O3(s) + 6Cl2(g) Compound

ΔGf°, kJ/mol

FeCl3(s)

–334.0

O2(g)

0

Fe2O3(s)

–742.2

Cl2(g)

0

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Chapter 17 Thermodynamics a. –148.4 kJ; spontaneous b. +148.4 kJ; nonspontaneous c. –408.2 kJ; spontaneous d. +408.2 kJ; nonspontaneous e. –375.0 kJ; spontaneous ANSWER: a 58. What is ΔG given ΔG° = 27.3 kJ/mol for the dissociation of PbCl2 at 25 °C if [Pb2+] and [Cl–] = 0.010 M? a. –6.93 kJ b. +59.3 kJ c. –4.28 kJ d. +7.98 kJ e. –11.2 kJ ANSWER: a 59. What is ΔG given ΔG° = –249.8 kJ/mol for the dissociation of AlCl3 at 25 °C if [Al3+] and [Cl–] = 0.35 M? AlCl3(s) ⇌ Al3+(aq) + 3Cl–(aq) a. –189.7 kJ b. –249.8 kJ c. –260.2 kJ d. +211.4 kJ e. +313.2 kJ ANSWER: c 60. For the reaction below, find ΔG at 298 K when PNO = 0.10 atm, P 2NO(g) + Cl2(g) → 2NOCl(g) a. –82.0 kJ b. –24.1 kJ c. –20.2 kJ d. +37.6 kJ e. +62.3 kJ ANSWER: b

= 0.25 atm, and PNOCl = 1.5 atm.

ΔG° = –41.0 kJ/mol

61. For the reaction below, find ΔG at 298 K when P

= 0.10 atm, P

= 0.25 atm, and P

= 1.5 atm.

N2(g) + 3H2(g) → 2NH3(g) Compound

°

ΔGf (kJ/mol)

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Chapter 17 Thermodynamics N2(g)

0

H2(g)

0

NH3(g) a. –16.4 kJ b. –32.8 kJ c. –14.8 kJ d. +40.9 kJ e. +11.3 kJ ANSWER: c

–16.4

62. Using the information below, determine the Gibbs free energy value associated with the production of 25.2 g of Fe2O3.

4Fe(s) + 3O2(g) → 2Fe2O3(s) Compound Fe(s) O2(g)

ΔGf°, kJ/mol 0 0

Fe2O3(s)

–742.2

ANSWER: –117 kJ 63. Using the information below, calculate the Gibbs free energy for the following reaction. Is the reaction spontaneous?

CO2(g) + 2HCl(g) → COCl2(g) + H2O(g) Compound

ΔGf°, kJ/mol

CO2(g) HCl(g) COCl2(g)

–394.4 –95.3 –204.9

H2O(g)

–228.6

ANSWER: 151.5 kJ The reaction is nonspontaneous.

64. Using the information below, calculate the Gibbs free energy for the following reaction. Is the reaction spontaneous?

N2O(g) + O3(g) → N2O4(g) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 17 Thermodynamics Compound

ΔGf°, kJ/mol

N2O(g)

103.7

O3(g)

163.2

99.8 N2O4(g) ANSWER: –167.1 kJ The reaction is spontaneous.

65. Which of the following statements is TRUE for a system at equilibrium? a. ΔG > 0 and K = Q b. ΔG < 0 and K = Q c. ΔG = 0 and K > Q d. ΔG = 0 and K < Q e. ΔG = 0 and K = Q ANSWER: e 66. Which of the following statements is TRUE for a system where Q < K? a. Gibbs free energy is positive and the reaction will proceed to the left. b. Gibbs free energy is negative and the reaction will proceed to the left. c. Gibbs free energy is positive and the reaction will proceed to the right. d. Gibbs free energy is negative and the reaction will proceed to the right. e. Gibbs free energy is zero and the reaction is at equilibrium.

ANSWER: d 67. Which of the following statements is TRUE for a system where Q > K? a. Gibbs free energy is positive and the reaction will proceed to the left. b. Gibbs free energy is negative and the reaction will proceed to the left. c. Gibbs free energy is positive and the reaction will proceed to the right. d. Gibbs free energy is negative and the reaction will proceed to the right. e. Gibbs free energy is zero and the reaction is at equilibrium.

ANSWER: a 68. Given the partial pressures of P

= 0.40 atm, P

= 0.35 atm, and P

= 0.60 atm, what is the

equilibrium constant, Kp, for the following reaction?

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Chapter 17 Thermodynamics 2SO2(g) + O2(g) → 2SO3(g) a. 1.26×1020

ΔG° = −142.0 kJ

b. 3.69×1024 c. 7.78×1024 d. 2.99×1022 e. 6.12×1023 ANSWER: c 69. What is the equilibrium constant, Kp, for the following equilibrium system?

2SO2(g) + O2(g) → 2SO3(g) a. 1.26×1020

ΔG° = −142.0 kJ

b. 3.69×1024 c. 7.78×1024 d. 2.99×1022 e. 6.12×1023 ANSWER: c 70. An unknown monoprotic acid (HX) has a Gibbs free energy value of 25.2 kJ for its ionization reaction.

What is the pKa of this acid?

HX(aq) ⇌ H+(aq) + X–(aq) a. 3.25 b. 4.42 c. 5.18 d. 6.20 e. 7.53 ANSWER: b

ΔG° = 25.2 kJ

71. What is the Ksp of calcium hydroxide?

Ca(OH)2(aq) ⇌ Ca2+(aq) + 2OH–(aq) a. 7.90×10–4

ΔG° = 29.5 kJ

b. 1.12×10–6 c. 3.68×10–6 d. 6.74×10–6

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Chapter 17 Thermodynamics e. 1.23×10–5 ANSWER: d 72. The partial pressures of each gas in the following system are PNO = 0.20 atm,

= 0.15 atm, and PNOCl =

0.80 atm. What are the Kp and Qp values for this system? At these pressures, will the reaction proceed to the right or to the left?

2NO(g) + Cl2(g) → 2NOCl(g) ANSWER Kp = 1.9 × 107 :

ΔG° = −40.9 kJ

Qp = 106.7 Kp is greater than Qp. Therefore, the reaction will proceed to the right to reach equilibrium.

73. Hydrogen phosphate has a Gibbs free energy value of 70.5kJ for its ionization reaction. What is the Ka and pKa of hydrogen phosphate based on ΔG° ?

HPO42–(aq) ⇌ H+(aq) + PO43–(aq) ANSWER: Ka = 4.4×10–13

ΔG° = 70.5 kJ

pKa = 12.36 74. The Ksp of an ionic compound was measured at two different temperatures. The data for this experiment is shown below. What is the ΔH° for the dissociation of the ionic compound?

Trial

Temperature (K)

Ksp

1

298

5.2×10–6

2

348

1.9×10–5

ANSWER: 22.3 kJ 75. The decomposition of dinitrogen tetroxide has a ΔH° of +58.0 kJ/mol and a Kp = 4.2×10−2 at 25 °C. What is the value of Kp at 75 °C? ANSWER: Kp = 1.2

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Chapter 16 Buffers and Titration 1. Which of the following best describes a buffer solution? a. A solution that completely neutralizes any added acid or base. b. A solution that resists changes in pH when small amounts of acid or base are added. c. A solution with a pH of 7.0. d. A solution that only contains a weak acid. e. A solution that only contains a weak base.

ANSWER: b 2. Which of the components in the following buffer solution would neutralize the addition of a small amount of a strong base?

HCO3–(aq) + H2O(l) ⇌H3O+(aq) + CO32–(aq) a. HCO3– b. H2O c. H3O+ d. CO32– e. Both HCO3– and CO32– ANSWER: a 3. Which of the components in the following buffer solution would neutralize the addition of a small amount of a strong acid?

H2PO4–(aq) + H2O(l) ⇌H3O+(aq) + HPO42–(aq) a. H2PO4– b. H2O c. H3O+ d. HPO42– e. Both H2PO4– and HPO42– ANSWER: d 4. Which of the following mixtures would make a good buffer solution when prepared as an aqueous solution? a. NaOH and HCl b. NaCl and HCl c. NaF and HF d. NaOH and HF e. HCl and HF ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration 5. Which of the following would cause the pH of a buffer solution to increase? a. adding a strong base b. adding a strong acid c. increasing the concentration of the weak acid component d. decreasing the concentration of the conjugate base component e. diluting the buffer solution with water ANSWER: a 6. Which of the following would cause the pH of a buffer solution to decrease? a. adding a strong base b. adding a strong acid c. decreasing the concentration of the weak acid component d. increasing the concentration of the conjugate base component e. diluting the buffer solution with water ANSWER: b 7. What happens to the pH of a buffer solution when a small amount of strong base is added? a. The pH increases significantly. b. The pH decreases significantly. c. The pH increases slightly. d. The pH decreases slightly. e. The pH remains unchanged.

ANSWER: c 8. What are the key components of a buffer solution?

ANSWER A buffer is a mixture of a weak acid and its conjugate base or a weak base and : its conjugate acid.

9. A buffer is made by dissolving 0.250 mol of both HC2H3O2 and NaC2H3O2 in 500.0 mL of water. Write out the reactions that would occur if (a) a small amount of a strong acid and is added to the buffer solution; (b) a small amount of a strong base is added to the buffer solution.

ANSWER: (a) H+(aq) + C2H3O2–(aq) → HC2H3O2(aq) (b) OH–(aq) + HC2H3O2(aq) → C2H3O2–(aq) + H2O(l) 10. A buffer is made by dissolving 0.100 mol of both NH3 and NH4Cl in 250.0 mL of water. Write out the reactions that would occur if (a) a small amount of a strong acid is added to the buffer solution; (b) a small amount of a strong base is added to the buffer solution.

ANSWER: (a) H+(aq) + NH3(aq) → NH4+(aq) (b) OH–(aq) + NH4+(aq) → NH3(aq) + H2O(l) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration 11. Describe the action of a buffer composed of a weak acid and its conjugate base when a strong acid or a strong base is added to the buffer.

ANSWER If a strong base is added to the buffer, the weak acid component of the buffer : reacts with it. If a strong acid is added to the buffer, the conjugate base component of the buffer reacts with it. The ratio of weak acid to conjugate base changes, but the amount of H+ and OH– present changes only slightly.

12. The Henderson-Hasselbalch equation is used to calculate a. the pH of any solution. b. the concentration of an acid solution. c. the pH of a weak acid solution. d. the pH of a buffer solution. e. the pH of a weak base solution.

ANSWER: d 13. If the concentration of the conjugate base in a buffer solution is increased, what will happen to the pH of the solution? a. The pH will increase. b. The pH will decrease. c. The pH will remain the same. d. The buffer capacity will decrease. e. The pKb of the conjugate base will increase.

ANSWER: a 14. A buffer solution is prepared with 0.20 M acetic acid (HC2H3O2) and 0.40 M sodium acetate (Na2H3O2).

The pKa of acetic acid is 4.76. What is the pH of this buffer solution? a. 4.46 b. 4.76 c. 5.06 d. 5.32 e. 9.24 ANSWER: c 15. A solution is prepared by dissolving 0.25 mol of acetic acid (HC2H3O2) and 0.15 mol of sodium acetate (NaC2H3O2) in water to a final volume of 400.0 mL. Calculate the pH of the buffer solution. The pKa of acetic acid is 4.76. a. 4.76 b. 4.24 c. 4.54 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration d. 5.12 e. 5.09 ANSWER: c 16. A buffer solution has a pH of 9.00. The pKa of the weak acid used to prepare the buffer is 8.50. What is the ratio of [base] to [acid] in this buffer solution? a. 0.32:1 b. 4.76:1 c. 3.16:1 d. 10.00:1 e. 0.10:1 ANSWER: c 17. Which of the following weak acid/conjugate base pairs would be most suitable for preparing a buffer solution with a pH of 5.0? a. formic acid/formate ion (pKa = 3.75) b. hydrofluoric acid/fluoride ion (pKa = 3.17) c. carbonic acid/bicarbonate ion (pKa = 6.37) d. ammonium ion/ammonia (pKa = 9.25) e. acetic acid/acetate ion (pKa = 4.76) ANSWER: e 18. A buffer solution is made using a weak acid with a Ka of 1.0×10–6. At what pH will the concentrations of the weak acid and its conjugate base be equal? a. 3.00 b. 6.00 c. 7.00 d. 9.00 e. 12.00 ANSWER: b 19. A buffer solution is prepared with 0.15 M benzoic acid and 0.20 M sodium benzoate. The Ka of benzoic acid is 6.3×10–5. What is the pH of this buffer solution? a. 5.63 b. 4.32 c. 6.21 d. 4.98 e. 9.27 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration ANSWER: b 20. What is the hydroxide concentration of a buffer solution that contains 0.25 M NH3 and 0.40 M NH4+? The Kb of ammonia (NH3) is 1.8×10–5. a. 3.6×10–10 M b. 4.2×10–5 M c. 1.7×10–4 M d. 1.1×10–5 M e. 9.8×10–4 M ANSWER: d 21. A buffer solution contains a weak acid and its conjugate base in a 1.5:1 ratio (acid:base). The pH of the buffer solution was measured to be 5.2. What is the pKa of the weak acid present in the buffer solution? a. 5.2 b. 6.2 c. 5.4 d. 7.6 e. 9.2 ANSWER: c 22. A buffer solution was prepared with 0.40 M HBrO and 0.30 M BrO– concentrations. The Ka of HBrO is 2.8×10–9. Set up an ICE table to determine the H+ ion concentration and then calculate the pH of this buffer solution.

ANSWER :

HBrO(aq)

⇌

H+(aq)

+

BrO–(aq)

Initial

0.40

-

0.30

Change

–x

+x

+x

[H+] = 3.7×10–9 M pH = 8.43 23. A buffer solution is prepared using HC2H3O2 and its conjugate base C2H3O2–. What does the concentration of HC2H3O2 need to be to make a buffer with a pH of 4.5 if [C2H3O2–] = 0.75 M? The pKa of HC2H3O2 is 4.76.

ANSWER: [HC2H3O2] = 1.4 M Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration 24. The pKa of HN3 is 4.72. What ratio of base to acid is needed to produce a buffer solution with pH of 5.0 with its conjugate base?

ANSWER:

25. Hypobromous acid, HBrO, has a pKa of 8.55. What is the most common form of hypobromous acid in an aqueous solution with a pH = 6.0?

ANSWER Since pH < pKa, the most common form of hypobromous acid in this solution : is the protonated form, HBrO. 26. A 1.0 L solution contains 0.20 mol NH3 and 0.40 mol NH4+. The Kb of ammonia (NH3) is 1.8×10–5. What are the concentrations of NH3 and NH4+ after 0.050 mol of HCl are added to the solution? Assume the additional HCl did not change the volume of the solution. a. [NH3] = 0.25 M and [NH4+] = 0.45 M b. [NH3] = 0.15 M and [NH4+] = 0.35 M c. [NH3] = 0.25 M and [NH4+] = 0.40 M d. [NH3] = 0.15 M and [NH4+] = 0.40 M e. [NH3] = 0.15 M and [NH4+] = 0.45 M ANSWER: e 27. A 1.00 L solution contains 1.25 mol HC2H3O2 and 1.50 mol C2H3O2–. The Ka of acetic acid (HC2H3O2) is 1.8×10–5. What are the concentrations of HC2H3O2 and C2H3O2– after 0.25 mol of KOH are added to the solution? Assume the additional KOH did not change the volume of the solution. a. [HC2H3O2] = 1.50 M and [C2H3O2–] = 1.75 M b. [HC2H3O2] = 1.00 M and [C2H3O2–] = 1.25 M c. [HC2H3O2] = 1.50 M and [C2H3O2–] = 1.50 M d. [HC2H3O2] = 1.00 M and [C2H3O2–] = 1.75 M e. [HC2H3O2] = 1.50 M and [C2H3O2–] = 1.25 M ANSWER: d 28. A 1.00 L solution contains 2.50 mol HC3H5O2 and 2.00 mol C3H5O2–. The Ka of propanoic acid (HC3H5O2) is 1.35×10–5. What are the concentrations of HC3H5O2 and C3H5O2– after 0.50 mol of KOH are added to the solution? a. [HC3H5O2] = 2.00 M and [C3H5O2–] = 2.50 M b. [HC3H5O2] = 2.00 M and [C3H5O2–] = 1.50 M c. [HC3H5O2] = 3.00 M and [C3H5O2–] = 2.50 M d. [HC3H5O2] = 3.00 M and [C3H5O2–] = 2.00 M Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration e. [HC3H5O2] = 2.25 M and [C3H5O2–] = 1.75 M ANSWER: a 29. A buffer solution is prepared by mixing 0.300 mol of formic acid (HCO2H, pKa = 3.75) with 0.450 mol of its conjugate base, formate (HCO2–). If 0.050 mol of NaOH is added to the buffer, what is the new pH of the solution? The final volume of the solution is 1.00 L. a. 3.75 b. 4.05 c. 3.00 d. 11.00 e. 5.10 ANSWER: b 30. A 2.00 L buffer solution contains 0.250 mol of ammonia (NH3, pKb = 4.75) and 0.400 mol of ammonium chloride (NH4Cl). Calculate the pH after the addition of 0.030 mol of HCl to the buffer. Assume the volume did not change after the addition of the HCl. a. 8.95 b. 4.75 c. 9.54 d. 4.46 e. 6.72 ANSWER: a 31. What is the change in pH after 0.080 mol of KOH is added to a buffer created by dissolving 0.550 mol of acetic acid (HC2H3O2, pKa = 4.76) and 0.600 mol of sodium acetate (NaC2H3O2)? The volume of the solution is 2.500 L of solution. a. 0.32 b. 1.06 c. 0.54 d. 0.12 e. 0.08 ANSWER: d 32. You have prepared a buffer solution by mixing 0.400 mol of hypochlorous acid (HClO, Ka = 3.5×10–8) with 0.250 mol of sodium hypochlorite (NaClO) in a total volume of 1.50 L. (a) Calculate the initial pH of this buffer solution. (b) Determine the pH after adding 0.020 mol of solid NaOH to the buffer.

ANSWER: (a) pH = 7.25 (b) pH = 7.31 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration 33. A buffer solution is made by combining 0.350 mol of benzoic acid (HC7H5O2, Ka = 6.3×10–5) and 0.500 mol of sodium benzoate (NaC7H5O2) in enough water to create a 2.00 L solution. (a) What is the pH of the original buffer solution? (b) What will the pH be after the addition of 0.040 mol of HCl? ANSWER: (a) pH = 4.36 (b) pH = 4.27 34. A 1.00 L buffer solution initially contains 0.150 M of hydrofluoric acid (HF, pKa = 3.20) and 0.200 M of sodium fluoride (NaF). (a) What is the change in pH after adding 0.015 mol of KOH to the buffer? (b) What would the change in pH be if the same amount of KOH (0.015 mol) was added to 1.00 L of pure water instead?

ANSWER: (a) The change in pH after adding KOH to the buffer is 0.08. (b) The change in pH after adding KOH to pure water is 5.18.

35. Which of the following best describes the buffer capacity of a solution? a. The pH range over which a buffer is effective. b. The concentration of the weak acid in the buffer solution. c. The ability to resist changes in pH upon the addition of strong acid or base. d. The ratio of weak acid to weak base in the buffer solution. e. The volume of the buffer solution.

ANSWER: c 36. Which of the following is the ideal choice for a buffer solution when the target pH is 7.0? a. a weak acid with a pKa of 5.2 b. a weak acid with a pKa of 7.3 c. a weak acid with a pKa of 9.0 d. a weak acid with a pKa of 3.2 e. a weak base with a pKb of 4.3 ANSWER: b 37. What is the effective range of a buffer solution? a. pH values within ±0.1 units of the pKa b. pH values within ±1.0 units of the pKa c. pH values within ±2.0 units of the pKa d. pH values less than 7.0 e. pH values equal to the pKa ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration 38. Which of the following describes a valid way to prepare a buffer solution using a pH meter? a. Mix calculated amounts of a weak acid and its conjugate base in the correct ratio. b. Completely neutralize a weak acid with a strong base while monitoring the pH. c. Add a strong acid to a weak base solution until the desired pH is reached. d. Dissolve the salt of a weak acid in water and measure the pH. e. Mix a strong acid and a strong base in equal proportions.

ANSWER: c 39. What ratio of HPO42– to H2PO4– is required to produce a buffer with a pH of 7.05? The pKa2 for phosphoric acid is 7.21. a. 0.69 b. 1.2 c. 1.5 d. 1.1 e. 0.73 ANSWER: a 40. What ratio of formate (HCO2–) to formic acid (HCO2H) is necessary to prepare a buffer solution with a pH of 4.00? The Ka of formic acid is 1.8×10–4. a. 0.75 b. 1.8 c. 1.2 d. 0.65 e. 1.3 ANSWER: b 41. Methylamine (CH3NH2) is a weak base with a pKb of 3.34. A buffer solution is prepared by mixing methylamine and its conjugate acid, CH3NH3+. If the desired pH of the buffer is 10.50, what should be the ratio of [CH3NH2] to [CH3NH3+] in the solution? a. 0.45 b. 1.9 c. 1.3 d. 0.89 e. 0.69 ANSWER: e 42. What mass of sodium acetate (NaC2H3O2) is needed to prepare 500.0 mL of an acetate buffer where the Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration total concentration of acetic acid (HC2H3O2) and acetate ion (C2H3O2–) is 1.50 M? a. 52.3 g b. 18.9 g c. 30.8 g d. 61.5 g e. 75.2 g ANSWER: d 43. A bicarbonate buffer system utilizes a mixture of carbonic acid (H2CO3) and bicarbonate ion (HCO3–). If you need to make 250.0 mL of this buffer with a total carbonate concentration of 0.400 M, what mass of sodium bicarbonate (NaHCO3) should be used? a. 8.40 g b. 16.8 g c. 4.20 g d. 33.6 g e. 17.1 g ANSWER: a 44. What ratio of benzoate (C7H5O2–) to benzoic acid (HC7H5O2) is necessary to prepare a buffer solution with a pH of 4.75? The pKa of benzoic acid is 4.20.

ANSWER: 3.55 45. What should the ratio of [NH3] to [NH4+] be for a buffer with a pH of 9.50? Ammonia, NH3, is a weak base with a pKb of 4.75.

ANSWER: 1.78 46. What mass of sodium chlorite (NaClO2) is needed to prepare 750.0 mL of a chlorite buffer solution with the total concentration of chlorous acid (HClO2) and chlorite ion (ClO2–) at 2.00 M? ANSWER: 136 g 47. Provide step-by-step instructions on preparing 500.0 mL of a phosphate buffer solution with a pH of 7.20 and a total buffer concentration (H2PO4– and HPO42–) of 0.500 M. The solution will be prepared using a pH meter, solid potassium dihydrogen phosphate (KH2PO4), and aqueous sodium hydroxide (NaOH). Include the mass of KH2PO4 needed and an outline for the procedure.

ANSWER 1. Measure out 34.0 g of KH2PO4 on a balance. : 2. Dissolve the KH2PO4 in approximately 450 mL of water.

3. Add the NaOH solution in small increments, while measuring the pH of the solution with a pH meter, until the pH reaches 7.20.

4. Transfer the solution to a volumetric flask and dilute to the final volume of 500.0 mL. Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration 48. Which of the following statements is correct regarding the titration of a strong base with a strong acid? a. The pH at the equivalence point will be equal to 7. b. The pH at the equivalence point will be greater than 7. c. The pH at the equivalence point will be less than 7. d. The pH at the equivalence point will depend on the concentration of the base. e. The pH at the equivalence point will depend on the concentration of the acid.

ANSWER: a 49. Which of the following statements is correct regarding the titration of a weak base with a strong acid? a. The pH at the equivalence point will be equal to 7. b. The pH at the equivalence point will be greater than 7. c. The pH at the equivalence point will be less than 7. d. The pH at the equivalence point will depend on the concentration of the base. e. The pH at the equivalence point will depend on the concentration of the acid.

ANSWER: c 50. Which of the following statements is correct regarding the titration of a weak acid with a strong base? a. The pH at the equivalence point will be equal to 7. b. The pH at the equivalence point will be greater than 7. c. The pH at the equivalence point will be less than 7. d. The pH at the equivalence point will depend on the concentration of the base. e. The pH at the equivalence point will depend on the concentration of the acid.

ANSWER: b 51. What type of titration would result in the following titration curve?

a. A weak acid being titrated with a strong base.

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Chapter 16 Buffers and Titration b. A strong acid being titrated with a strong base. c. A weak base being titrated with a strong acid. d. A diprotic weak acid being titrated with a strong base. e. A diprotic strong acid being titrated with a strong base.

ANSWER: c 52. What type of titration would result in the following titration curve?

a. A weak acid being titrated with a strong base. b. A strong acid being titrated with a strong base. c. A weak base being titrated with a strong acid. d. A diprotic weak acid being titrated with a strong base. e. A dibasic weak base being titrated with a strong acid.

ANSWER: d 53. A 25.0 mL sample of 0.100 M acetic acid (HC2H3O2, Ka = 1.8×10–5) is titrated with 0.100 M NaOH. What is the pH at the equivalence point? a. 2.87 b. 4.74 c. 7.86 d. 8.72 e. 9.13 ANSWER: d 54. A student titrates 20.0 mL of an unknown, monoprotic weak acid with 0.150 M NaOH. The titration curve shows an equivalence point at 28.0 mL of added NaOH. What is the original concentration of the weak acid? a. 0.0705 M Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration b. 0.105 M c. 0.210 M d. 0.150 M e. 0.125 M ANSWER: c 55. A 25.0 mL sample of 0.100 M methylamine (CH3NH2, Kb = 4.4×10–4) is titrated with 0.100 M HCl. What is the pH at the equivalence point? a. 2.92 b. 4.78 c. 5.97 d. 7.14 e. 10.12 ANSWER: c 56. The titration of a weak acid by NaOH produced the following titration curve. Estimate the pKa of the acid using the titration curve.

a. 3.5 b. 4.8 c. 6.2 d. 7.9 e. 12.2 ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 16 Buffers and Titration 57. The titration of a weak base by HCl produced the following titration curve. Estimate the pKb of the base using the titration curve.

a. 10.5 b. 9.2 c. 7.8 d. 6.1 e. 1.8 ANSWER: b 58. When titrating a weak base with a strong acid, at what point is the pH equal to the pKb of the weak base? a. At the start of the titration b. At the equivalence point c. At the half-equivalence point d. When the concentration of the weak base is twice the concentration of its conjugate acid e. When the concentration of the weak base is half the concentration of its conjugate acid ANSWER: c 59. Sketch a titration curve between a weak acid and a strong base. Label each axis, the equivalence point, and the half-equivalence point on the curve.

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Chapter 16 Buffers and Titration ANSWER:

60. Sketch a titration curve between a weak base and a strong acid. Label each axis, the equivalence point, and the half-equivalence point on the curve.

ANSWER:

61. A 15.0 mL monoprotic weak acid was titrated with a 0.155 M NaOH hydroxide solution. The titration produced the data and titration curve below. Using that information, what is the original concentration and the pKa of the weak acid?

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Chapter 16 Buffers and Titration

ANSWER: [HA] = 0.258 M; pKa = 5.63 62. What is the pH of the solution at the point in a titration when 7.80 mL of 0.255 M NaOH is added to a 20.00 mL 0.175 M HCl solution?

ANSWER: pH = 1.26 63. What is the pH of the solution at the point in a titration when 11.80 mL of 0.210 M NaOH is added to 15.00 mL of a 0.185 M HC2H3O2 solution? The Ka of HC2H3O2 is 1.8×10–5.

ANSWER: pH = 5.66 64. What is the pH of the solution at the point in a titration when 14.80 mL of a 0.155 M HCl solution is added to 20.00 mL of a 0.175 M NH3 solution? The Kb of NH3 is 1.8×10–5.

ANSWER: pH = 8.98 65. Which of the following statements is NOT consistent with the Lewis model of acids and bases? a. A Lewis acid is an electron pair acceptor. b. A Lewis base is an electron pair donor. c. A Lewis acid-base reaction involves the transfer of a proton. d. A Lewis acid-base reaction results in the formation of a covalent bond.

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Chapter 16 Buffers and Titration e. All of the above statements are consistent with the Lewis model.

ANSWER: c 66. Which of the following is an example of a Lewis acid but NOT a Brønsted-Lowry acid? a. H2O b. HCl c. NH3 d. BF3 e. HC2H3O2 ANSWER: d 67. Which of the following can react as a Lewis acid? a. O2– b. Fe3+ c. NH3 d. F– e. OH– ANSWER: b 68. In the formation of a coordination compound, the central metal ion acts as a a. Lewis base by donating an electron pair. b. Lewis acid by accepting an electron pair. c. Arrhenius acid by producing an H+ ion. d. Brønsted-Lowry acid by donating a proton. e. Brønsted-Lowry base by accepting a proton.

ANSWER: b 69. Which of the following best describes a coordination complex? a. A compound composed solely of hydrogen bonds. b. A compound formed by ionic bonding between metal ions and nonmetals. c. A compound where metal ions are bonded to several ligands. d. A compound that minimized changes in pH. e. A compound where metal ions donate electrons to organic ligands.

ANSWER: c 70. Which of the following statements is true regarding the following reaction?

2NH3 + Ag+ → [H3N–Ag–NH3]+ a. NH3 is acting as a Lewis base.

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Chapter 16 Buffers and Titration b. NH3 is acting as a Brønsted-Lowry base. c. Ag+ is acting as a Lewis base. d. [H3N–Ag–NH3]+ is acting as a Lewis base. e. [H3N–Ag–NH3]+ is acting as a Lewis acid.

ANSWER: a 71. Which of the following statements is true regarding the following reaction?

O2– + SO3 → SO42– a. O2– is acting as a Lewis acid. b. O2– is acting as a Brønsted-Lowry base. c. SO3 is acting as a Lewis acid. d. SO42– is acting as a Lewis base. e. SO42– is acting as a Lewis acid.

ANSWER: c 72. What is the definition of an acid using the Lewis acid/base model?

ANSWER According to the Lewis acid/base model, an acid is a substance that can accept : a pair of electrons to form a new covalent bond.

73. What is the definition of a base using the Lewis acid/base model?

ANSWER: A Lewis base is a species that donates an electron pair to form a new covalent bond.

74. What species in the following reaction would be classified as a Lewis acid?

BF3 + NH3 → F3B–NH3 ANSWER: BF3 75. What species in the following reaction would be classified as a Lewis base?

BF3 + NH3 → F3B–NH3 ANSWER: NH3

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Chapter 15 Acids and Bases 1. According to the Arrhenius definition, what substance do acids produce in water? a. H+ b. OH– c. H2 d. O2 e. H2O ANSWER: a 2. According to the Arrhenius definition, what substance do bases produce in water? a. H+ b. OH– c. H2 d. O2 e. H2O ANSWER: b 3. Which of the following is an example of a polyprotic acid? a. HCl b. HC2H3O2 c. CH3OH d. NH3 e. H3PO4 ANSWER: e 4. Which substance is the acid in the following reaction?

HClO4(aq) + CH2NH2(aq) → ClO4–(aq) + CH2NH3+(aq) a. HClO4 b. CH2NH2 c. ClO4– d. CH2NH3+ e. No substance in the reaction is an acid.

ANSWER: a 5. Which substance is the base in the following reaction?

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Chapter 15 Acids and Bases HI(aq) + C2H5NH2(aq) → I–(aq) + C2H5NH3+(aq) a. HI b. C2H5NH2 c. I– d. C2H5NH3+ e. No substance in the reaction is a base.

ANSWER: b 6. What is the term for a substance that can release more than one H+ ion into aqueous solution? a. Arrhenius acid b. Arrhenius base c. polyprotic acid d. deprotonation e. protonation ANSWER: c 7. What is the Brønsted-Lowry definition for acids?

ANSWER: Acids are compounds that donate H+ ions, often called proton donors.

8. What is the Brønsted-Lowry definition for bases?

ANSWER: Bass are compounds that accept H+ ions, often called proton acceptors.

9. Write two balanced chemical equations showing the acidic properties of hydrobromic acid, HBr, one equation with water not included and the second with water included.

ANSWER: First equation: HBr(aq) → H+(aq) + Br–(aq) Second equation: HBr(aq) + H2O(l) → H3O+(aq) + Br–(aq) 10. Write the balanced chemical equation that shows the basic properties of solid lithium hydroxide, LiOH, when it is dissolved in water.

ANSWER: LiOH(s) → Li+(aq) + OH–(aq) 11. Which of the following weak acids is the strongest? a. H2CO3, Ka = 4.5 × 10–7 b. HN3, Ka = 1.9 × 10–5 c. HF, Ka = 6.3 × 10–4 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases d. HBrO, Ka = 2.8 × 10–9 e. H2S, Ka = 8.9 × 10–8 ANSWER: c 12. Which of the following weak acids is the weakest? a. H2CO3, Ka = 4.5 × 10–7 b. HN3, Ka = 1.9 × 10–5 c. HF, Ka = 6.3 × 10–4 d. HBrO, Ka = 2.8 × 10–9 e. H2S, Ka = 8.9 × 10–8 ANSWER: d 13. Which of the following weak bases is the strongest? a. NH3, Kb = 1.8 × 10–5 b. CO32–, Kb = 2.1 × 10–4 c. PO43–, Kb = 2.1× 10–2 d. (C2H5)3N, Kb = 5.6 × 10–4 e. CH5N, Kb = 4.6 × 10–4 ANSWER: c 14. Which of the following is the conjugate acid of HSO3–? a. SO32– b. H2SO3 c. H3O+ d. HSO4– e. SO42– ANSWER: b 15. Which of the following are conjugate acid-base pairs involved in the following reaction?

H2PO4–(aq) + H2O(l) ⇌H3O+(aq) + HPO42–(aq) a. H2PO4– and H2O b. H3O+ and HPO42– c. H2PO4– and H3O+ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases d. H2PO4– and HPO42– e. H2O and HPO42– ANSWER: d 16. Write the equation for the equilibrium reaction that the weak base methyl amine, CH3NH2, would undergo in an aqueous solution. Indicate each conjugate acid-base pair that is part of the reaction.

ANSWER: CH3NH2(aq) + H2O(l) ⇌CH3NH3+(aq) + OH–(aq) Base: CH3NH2 Conjugate acid: CH3NH3+ Acid: H2O Conjugate base: OH– 17. Write the equation for the equilibrium reaction that the weak acid hypochlorous acid, HClO, would undergo in an aqueous solution. Indicate each conjugate acid-base pair that is part of the reaction.

ANSWER: HClO(aq) + H2O(l) ⇌H3O+(aq) + ClO–(aq) Acid: HClO Conjugate base: ClO– Base: H2O Conjugate acid: H3O+ 18. Which of the following is an example of a neutralization reaction? a. HClO4(aq) + KOH(aq) → KClO4(aq) + H2O(l) b. 2Al(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2(g) c. Mg(s) + 2H+(aq) → Mg2+(aq) + H2(g) d. CO2(g) + H2O(l) ⇌H2CO3(aq) e. H2SO4(aq) + H2O(l) → H3O+(aq) + HSO4–(aq) ANSWER: a 19. Which of the following is an example of a neutralization reaction? a. Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g) b. 2Sc(s) + 6HI(aq) → 2ScI3(aq) + 3H2(g) c. SO3(g) + H2O(l) → H2SO4(aq) d. HCO3–(g) + H2O(l) ⇌H2CO3(aq) + OH–(aq) e. HCl(aq) + LiOH(aq) → LlCl(aq) + H2O(l) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases ANSWER: e 20. Which of the following is an example of a reaction between an acid and a metal? a. SO2(g) + H2O(l) ⇌H2SO3(aq) b. Sn(s) + 2H+(aq) → Sn2+(aq) + H2(g) c. Ag+(aq) + Cl–(aq) ⇌AgCl(s) d. Al3+(aq) + 4OH–(aq) ⇌Al(OH)4–(aq) e. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) ANSWER: b 21. Which of the following is an example of a reaction between an acid and a metal? a. 2NO2(g) + H2O(l) → HNO2(aq) + HNO3(aq) b. Pb2+(aq) + Cl–(aq) ⇌PbCl2(s) c. Cu(s) + 2HNO3(aq) → Cu(NO3)2(aq) + H2(g) d. Ag+(aq) + 2S2O32–(aq) ⇌Ag(S2O3)23–(aq) e. H+(aq) + OH–(aq) ⇌H2O(l) ANSWER: c 22. Which of the following is an example of an acid formation reaction from nonmetal oxides? a. 2H2O(l) ⇌H3O+(aq) + OH–(aq) b. Pb(s) + 2H+(aq) → Pb2+(aq) + H2(g) c. Cu(s) + 2 HNO3(aq) → Cu(NO3)2(aq) + H2(g) d. 2NO2(g) + H2O(l) → HNO2(aq) + HNO3(aq) e. H2CrO4(aq) → 2H+(aq) + CrO42–(aq) ANSWER: d 23. Which of the following is an example of an acid formation reaction from nonmetal oxides? a. CaCO3(s) ⇌CaO(s) + CO2(g) b. Ca(s) + 2H+(aq) → Ca2+(aq) + H2(g) c. H2CO3(aq) ⇌H+(aq) + HCO3–(aq) d. HBrO4(aq) ⇌H+(aq) + BrO4–(aq) e. Cl2O(g) + H2O(l) → 2HClO(aq) ANSWER: e 24. Write a balanced chemical equation for the neutralization of aqueous nitric acid with aqueous strontium hydroxide.

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Chapter 15 Acids and Bases ANSWER: 2HNO3(aq) + Sr(OH)2(aq) → Sr(NO3)2(aq) + 2H2O(l) 25. Write a balanced chemical equation for the reaction between solid diphosphorus pentoxide (P2O5) and water to form phosphoric acid.

ANSWER: P2O5(s) + 3H2O(l) ⇌2H3PO4(aq) 26. Write a balanced chemical equation for the reaction between chromium metal and nitric acid to form aqueous chromium(III) nitrate and hydrogen gas.

ANSWER: 2Cr(s) + 6HNO3(aq) → 2Cr(NO3)3(aq) + 3H2(g) 27. Which of the following equations represents the autoionization of water? a. H+(aq) + H2O(l) ⇌H3O+(aq) b. 2H2O(l) ⇌H3O+(aq) + OH–(aq) c. H2O(l) + OH–(aq) ⇌H3O2–(aq) d. H3O+(aq) ⇌2H+(aq) + OH–(aq) e. H+(aq) + OH–(aq) ⇌H2O(l) ANSWER: b 28. The concentration of H+ in an aqueous solution was measured to be 1.25×10–9 M. What is the concentration of OH– present in this solution? Is the solution acidic, basic, or neutral? a. The concentration of OH– is 8.0×10–6 M. The solution is acidic. b. The concentration of OH– is 2.5×10–4 M. The solution is acidic. c. The concentration of OH– is 8.0×10–6 M. The solution is basic. d. The concentration of OH– is 2.5×10–4 M. The solution is basic. e. The concentration of OH– is 1.5×10–3 M. The solution is acidic.

ANSWER: c 29. The concentration of OH– in an aqueous solution was measured to be 3.5×10–4 M. What is the concentration of H+ present in this solution? Is the solution acidic, basic, or neutral? a. The concentration of H+ is 4.7×10–9 M. The solution is acidic. b. The concentration of H+ is 2.9×10–11 M. The solution is acidic. c. The concentration of H+ is 4.7×10–9 M. The solution is basic. d. The concentration of H+ is 2.9×10–11 M. The solution is basic. e. The concentration of H+ is 3.2×10–6 M. The solution is acidic.

ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases 30. Which of the following is true for a neutral solution? a. [OH–] > [H3O+] b. [OH–] < [H3O+] c. [H3O+] = 0 M d. [H3O+] = [OH–] = 0 M e. [H3O+] = [OH–] ANSWER: e 31. What is the pH of a solution that has a [H+] = 0.0020 M? a. 2.00 b. 2.70 c. 3.10 d. 4.00 e. 4.30 ANSWER: b 32. What is the pH of a solution that has a [OH–] = 0.0015 M? a. 2.82 b. 4.26 c. 9.98 d. 11.18 e. 12.52 ANSWER: d 33. What is the [H+] for a solution that has a pH of 4.25? a. 4.2×10–5 M b. 5.6×10–5 M c. 3.9×10–4 M d. 2.7×10–4 M e. 3.2×10–6 M ANSWER: b 34. What is the pH of a solution that has a [H+] = 4.3×10–5 M? Is the solution acidic, basic, or neutral?

ANSWER: The pH is 4.37; the solution is acidic.

35. What is the pH of a solution that has a [OH–] = 7.2×10–9 M? Is the solution acidic, basic, or neutral?

ANSWER: The pH is 5.86; the solution is acidic.

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Chapter 15 Acids and Bases 36. The pH of human blood is approximately 7.40. What is the concentration of OH– and H+ present in blood?

Is blood solution acidic, basic, or neutral?

ANSWER The concentration of H+ in blood is 4.0 × 10–8 M, and the concentration of : OH– is 2.5 × 10–7 M. Blood is slightly basic.

37. What is the pH of 0.125 M HCl solution? a. 1.239 b. 2.040 c. 0.903 d. 1.754 e. 3.210 ANSWER: c 38. What is the concentration of a HNO3 solution that has a pH of 2.50? a. 0.019 M b. 2.4×10–3 M c. 0.25 M d. 0.12 M e. 3.2×10–3 M ANSWER: e 39. What is the pH of a 0.50 M HCN solution? The Ka of HCN is 6.2×10–10. a. 1.25 b. 4.75 c. 6.31 d. 3.92 e. 2.25 ANSWER: b 40. What is the percent ionization of a 0.75 M HF solution? The Ka of HF is 6.3×10–4. a. 1.9% b. 3.6% c. 6.3% d. 2.9% e. 4.1% ANSWER: d 41. What is the pH of a 0.075 M NaOH solution? a. 1.12 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases b. 4.25 c. 7.00 d. 8.75 e. 12.88 ANSWER: e 42. What is the pH of a 0.00053 M KOH solution? a. 2.28 b. 10.72 c. 7.00 d. 11.72 e. 3.28 ANSWER: b 43. What is the pH of a 0.0035 M Ba(OH)2 solution? a. 2.15 b. 2.45 c. 12.75 d. 11.85 e. 11.55 ANSWER: d 44. What is the pH of a 0.030 M CH3NH2 solution? The Kb of CH3NH2 is 4.6×10–4. a. 2.43 b. 9.12 c. 11.57 d. 8.90 e. 12.07 ANSWER: c 45. What is the pH of a 0.025 M of phosphoric acid? The Ka values for phosphoric acid are shown.

H3PO4 Ka1 = 6.9×10–3 H2PO4– Ka2 = 6.2×10–8 HPO42– Ka3 = 4.8×10–12 a. 1.25 b. 1.88 c. 3.28 d. 4.26 e. 5.13 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases ANSWER: b 46. What is the concentration of H+ ions present in a 1.05 M HN3 solution? The Ka of HN3 is 1.9×10–5. a. 4.5×10–3 M b. 2.7×10–3 M c. 1.3×10–3 M d. 9.8×10–4 M e. 5.6×10–4 M ANSWER: a 47. What is the concentration of H+ ions and the pH of a 1.25 M acetic acid solution? The Ka of acetic acid, HC2H3O2, is 1.8×10–5.

ANSWER The concentration of [H+] in the solution is 4.7×10–3 M. The pH of the solution : is 2.32.

48. What is the pH of a solution in which 1.25 g of NaOH was dissolved in enough water to make a 250.0 mL solution?

ANSWER: The pH of the solution is 13.100 49. A solution containing an unknown monoprotic acid with a concentration of 0.100 M has a pH of 4.23. What is the Ka of this unknown acid?

ANSWER: The Ka of the unknown acid is 3.46×10−8.

50. Potassium hydrogen phthalate, often called KHP, is a weak acid commonly used to determine the concentration of bases. What is the pH of a KHP solution made by dissolving 0.500 g of KHP in 50.0 mL of solution? The Ka of KHP is 4.0×10–6, and the molar mass of KHP is 204.22 g/mol.

ANSWER: The pH of the solution is 3.35.

51. What is the concentration of OH– ions and the pH of a 1.5 M pyridine solution? The Kb of pyridine, C5H5N, is 1.6×10–9.

ANSWER: The concentration of [OH−] is 4.9×10−5 M. The pH of the solution is 9.69.

52. What is the equation for the basic reaction of the conjugate base of acetic acid (HC2O3O2) that would occur in an aqueous solution? a. HC2O3O2(aq) + OH–(aq) ⇌H2O(l) + C2H3O2–(aq) b. HC2O3O2(aq) + NH3(aq) ⇌NH4+(aq) + C2H3O2–(aq) c. C2H3O2–(aq) + H3O+(aq) ⇌HC2H3O2(aq) + H2O(l) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases d. C2H3O2–(aq) + H2O(l) ⇌C2H3O2–(aq) + OH–(aq) e. C2H3O2–(aq) + OH–(aq) + 2H+(aq) ⇌HC2H3O2(aq) + H2O(l) ANSWER: d 53. What is the equation for the acidic reaction of the conjugate acid of ammonia (NH3) that would occur in an aqueous solution? a. NH4+(aq) + OH–(aq) ⇌H2O(l) + NH3(aq) b. H2O(l) + NH3(aq) ⇌NH4+(aq) + OH–(aq) c. NH4+(aq) + H2O(l) ⇌NH3(aq) + H3O+(aq) d. NH4+(aq) ⇌NH3(aq) + H+(aq) e. NH4+(aq) + H3O+(aq) + 2OH–(aq) ⇌NH3(aq) + 3H2O(l) ANSWER: c 54. The Ka of HBrO is 2.8×10–9. What is the Kb of its conjugate base, BrO–? a. 4.4×10–4 b. 1.9×10–7 c. 6.2×10–8 d. 1.5×10–6 e. 3.6×10–6 ANSWER: e 55. The Kb of PO43– is 2.1×10–2. What is the Ka of its conjugate acid, HPO42–? a. 9.4×10–5 b. 4.8×10–13 c. 1.2×10–8 d. 7.2×10–11 e. 4.1×10–12 ANSWER: b 56. What is the pKa of an acid that has a Ka of 1.8×10–7. a. 1.87 b. 4.83 c. 6.74 d. 5.49 e. 9.81 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases ANSWER: c 57. What is the pKb of a base that has a Kb value of 2.3×10–5. a. 1.32 b. 4.64 c. 7.25 d. 2.35 e. 11.28 ANSWER: b 58. The amino acid cysteine has a pKa value of 1.71. What is the associated Ka value for cysteine? a. 1.9×10–2 b. 5.6×10–2 c. 3.3×10–3 d. 4.7×10–3 e. 1.3×10–2 ANSWER: a 59. Classify each compound as either acidic, basic, or neutral when dissolved in water: NaF, NaNO3, KCl, K2CO3, NH4Cl. ANSWER: NaF - basic NaNO3 - neutral KCl - neutral K2CO3 - basic NH4Cl - acidic 60. Identify which of the following compounds would produce a basic solution when dissolved in water: NaF, NaCl, NaBr, NaI. ANSWER NaF would produce a basic solution when dissolved in water. The other : compounds would produce a neutral solution.

61. The hydrogen carbonate ion, HCO3–, can react like an acid and as a base. Write the two chemical equations illustrating both these reactions.

ANSWER: Acidic reaction: HCO3–(aq) + H2O(l) ⇌H3O+(aq) + CO32–(aq) Basic reaction: HCO3–(aq) + H2O(l) ⇌OH–(aq) + H2CO3(aq) 62. Which of the following trends in the periodic table best explains the increase in acidity of hydrogen halides (HX) as you move down the group?

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Chapter 15 Acids and Bases a. Increasing electronegativity of halogens b. Increasing atomic size c. Increasing bond dissociation energy d. Increasing bond strength e. Decreasing electronegativity of hydrogen ANSWER: b 63. How does resonance affect the strength of acids? a. It makes acids weaker by destabilizing the conjugate base. b. It makes acids stronger by stabilizing the conjugate base. c. It does not affect acid strength. d. It makes acids stronger by destabilizing the conjugate acid. e. It makes acids weaker by stabilizing the conjugate acid.

ANSWER: b 64. Which concept explains the increased acidity of trichloroacetic acid (CCl3COOH) compared to acetic acid (CH3COOH)? a. resonance b. steric hindrance c. electron affinity d. inductive effect e. atomic size ANSWER: d 65. Which of the following acids would you expect to be the weakest based on periodic trends? a. H2Se b. H2S c. HI d. HBr e. HCl ANSWER: b 66. Which of the following solutions would have the highest pH? a. 0.1 M HClO4 b. 0.1 M HClO3 c. 0.1 M HClO2 d. 0.1 M HClO Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases e. They would all have the same pH. ANSWER: d 67. Why is sulfuric acid, H2SO4, a stronger acid than sulfurous acid, H2SO3?

ANSWER Sulfuric acid is a stronger acid than sulfurous acid because it has more oxygen : atoms, which leads to greater induction and increased resonance stabilization of the conjugate base.

68. Rank the following substances in order of increasing acid strength: HBrO3, HBrO4, HBrO, HBrO2, BrO4–.

ANSWER: BrO4– < HBrO < HBrO2 < HBrO3 < HBrO4 69. Describe the periodic trend, both for going across a period and along a group, in terms acid strength.

ANSWER The strength of an acid increases both across a period (from left to right) and : down a group (from top to bottom) in the periodic table. Across a period, the increasing effective nuclear charge makes it easier to lose H+, increasing acidity. Down a group, the increasing size of the atom allows for better spreading of the negative charge on the conjugate base, also leading to increased acidity 70. In a titration experiment, 0.155 g of potassium hydrogen phthalate (KHP, 204.22 g/mol) was dissolved in water and used to standardize a sodium hydroxide solution. What is the concentration of the sodium hydroxide solution if it required 27.55 mL of the solution to titrate the KHP sample? KHP is monoprotic and reacts with NaOH in a one-to-one molar ratio. a. 0.0721 M b. 0.123 M c. 0.0981 M d. 0.0275 M e. 0.109 M ANSWER: d 71. Which statement best describes the use of litmus in acid-base reactions? a. Litmus is a synthetic indicator used only in alkaline solutions. b. Litmus is a natural indicator that changes color at a pH of 7. c. Litmus is a natural indicator that turns red in acidic solutions and blue in basic solutions. d. Litmus can only indicate the presence of strong acids. e. Litmus decreases the pH of the solution it is added to.

ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 15 Acids and Bases 72. What property of a pH indicator makes it useful for determining the end point of a titration? a. Its ability to change color at a specific pH range. b. Its ability to increase the pH of the solution. c. Its ability to decrease the pH of the solution d. Its ability to act as a catalyst in the reaction. e. Its solubility changes with pH changes.

ANSWER: a 73. What is the concentration of a 10.00 mL potassium hydroxide (KOH) solution that required 12.55 mL of a 0.125 M sulfuric acid (H2SO4) solution to reach the end point of a titration? a. 0.157 M b. 0.314 M c. 0.129 M d. 0.274 M e. 0.218 M ANSWER: b 74. In a titration experiment, 0.200 g of potassium hydrogen phthalate (KHP, 204.22 g/mol) is dissolved in water and used to standardize a potassium hydroxide (KOH) solution. The KOH solution has an approximate concentration of 0.15 M. How many mL of the KOH solution will be required to reach the end point of the titration? KHP is monoprotic and reacts with KOH in a one-to-one molar ratio.

ANSWER: 6.5 mL 75. What volume of a 2.5 M HCl solution is required to completely neutralize 1.5 L of a 0.050 M Ca(OH)2 solution?

ANSWER: 0.060 L

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Chapter 14 Equilibrium 1. Which statement best describes a system at dynamic equilibrium? a. The forward reaction rate is greater than the reverse reaction rate. b. The reverse reaction rate is greater than the forward reaction rate. c. The concentrations of reactants and products are equal. d. The reverse reaction rate is equal to the forward reaction rate. e. The forward and reverse reactions have stopped.

ANSWER: d 2. Which is the equilibrium constant expression for the following equilibrium system?

N2(g) + 3H2(g) ⇌2NH3(g) a. b.

c.

d. e.

ANSWER: b 3. Which is the equilibrium constant expression for the following equilibrium system?

Fe3+(aq) + SCN–(aq) ⇌FeSCN2+(aq) a. b. c. d.

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Chapter 14 Equilibrium e.

ANSWER: a 4. What is the corresponding chemical equation for the following equilibrium constant expression?

a. SO2(g) +

O2(g) ⇌SO3(g)

b.

SO2(g) + O2(g) ⇌ SO3(g)

c.

SO3(g) ⇌ SO2(g) + O2(g)

d. 2SO3(g) ⇌2SO2(g) + O2(g) e. 2SO2(g) + O2(g) ⇌2SO3(g) ANSWER: e 5. What is the corresponding chemical equation for the following equilibrium constant expression?

a.

O2(g) ⇌ O3(g)

b.

O3(g) ⇌ O2(g)

c. 2O2(g) ⇌3O3(g) d. 2O3(g) ⇌3O2(g) e. O32+(g) ⇌O23+(g) ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium 6. An equilibrium mixture of C4H8 and C2H4 has the following concentrations: [C4H8] = 0.0125 M and [C2H4] = 0.0450 M. What is the value of Kc for this mixture?

C4H8(g) ⇌2C2H4(g) a. 6.17 b. 0.162 c. 3.6 d. 0.278 e. 0.891 ANSWER: b 7. An equilibrium mixture of NO, Cl2, and NOCl has the following concentrations: [NO] = 0.0750 M, [Cl2] = 0.0155 M, and [NOCl] = 0.450 M. What is the value of Kc for this mixture?

2NO(g) + Cl2(g) ⇌2NOCl(g) a. 2.32×103 b. 387 c. 4.31×10–3 d. 0.691 e. 1.67×104 ANSWER: a 8. What is the equilibrium concentration of ammonia, NH3, if [H2] = 0.275 M and [N2] = 0.225 M? N2(g) + 3H2(g) ⇌2NH3(g)

Kc = 0.50

a. 0.15 b. 0.048 c. 0.18 d. 0.025 e. 2.3×10–3 ANSWER: b 9. What is the equilibrium concentration of H2 if [I2] = 0.075 M and [HI] = 0.62 M? H2(g) + I2(g) ⇌2HI(g)

Kc = 55

a. 0.15 b. 0.072 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium c. 0.093 d. 0.23 e. 0.015 ANSWER: c 10. What is the equilibrium constant expression for the following system?

2NH3(g) + 3I2(g) ⇌N2(g) + 6HI(g) ANSWER:

11. What is the equilibrium constant expression for the following system?

C2H2(g) + 2Br2(g) ⇌C2H2Br4(g) ANSWER:

12. What is the value of Kc for the following equilibrium system, given these equilibrium concentrations: [I2] = 8.69×10–3 M, [I–] = 7.23×10–3 M, and [I3–] = 5.55×10–3 M? I2(aq) + I–(aq) ⇌I3–(aq) ANSWER: 88.3 13. The energy diagram for a system at equilibrium is shown. Will this equilibrium favor the reactants or the products? Explain your answer.

ANSWER The equilibrium will favor the reactants because the reverse reaction has a : lower activation energy.

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Chapter 14 Equilibrium 14. Which is the equilibrium constant expression for the following equilibrium system?

HF(aq) + H2O(l) ⇌H3O+(aq) + F–(aq) a. b. c.

d. e.

ANSWER: c 15. Which species would NOT be included in the equilibrium constant expression for the following system?

Fe3+(aq) + 3C6H5COOH(aq) + 3H2O(l) ⇌3H3O+(aq) + Fe(C6H5COO)3(aq) a. Fe3+(aq) b. C6H5COOH(aq) c. H2O(l) d. H3O+(aq) e. Fe(C6H5COO)3(aq) ANSWER: c 16. Which species would NOT be included in the equilibrium constant expression for the following system?

BaSO4(s) + H3O+(aq) ⇌Ba2+(aq) + HSO4–(aq) + H2O(l) a. HSO4–(aq) b. Ba2+(aq) c. BaSO4(s) d. H3O+(aq) e. BaSO4(s) and H2O(l) ANSWER: e Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium 17. Which is the equilibrium constant expression for the following equilibrium system?

CaCO3(s) ⇌CaO(s) + CO2(g) a. b. c. d. e.

ANSWER: e 18. A 1.0-L flask contains 12.5 mg of MgCO3, 22.5 mg of MgO, and 129 mg of CO2. What is the value of Kc for this system?

MgCO3(s) ⇌MgO(s) + CO2(g) a. 1.11×10–2 b. 2.93×10–3 c. 4.19×10–2 d. 3.37×10–3 e. 6.38×10–4 ANSWER: b 19. Which is the solubility product expression for the dissolution of calcium phosphate, Ca3(PO4)2? a. b. c.

d.

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Chapter 14 Equilibrium e.

ANSWER: e 20. In a saturated solution of copper(I) chloride, CuCl, the concentrations of the ions are [Cu+] = 1.1×10–3 M and [Cl–] = 1.1×10–3 M. What is the Ksp for CuCl? a. 6.1×10–6 b. 2.3×10–6 c. 4.9×10–5 d. 1.2×10–6 e. 3.4×10–4 ANSWER: d 21. What is the concentration of Pb2+ ions in a saturated PbCl2 solution? The Ksp for PbCl2 is 1.6×10–5. a. 0.016 M b. 0.052 M c. 4.1×10–4 M d. 3.7×10–3 M e. 0.013 M ANSWER: a 22. What is the Kc expression for the following equilibrium system?

2KClO3(s) ⇌2KCl(s) + 3O2(g) ANSWER: 23. Chromium(III) hydroxide has a solubility of 520 mg/L in acidic conditions. What is the Ksp of Cr(OH)3 under these conditions?

ANSWER: 1.8×10–8 24. What is the value of Ksp expression for the dissolution of Ag3AsO4?

ANSWER: 25. The Ksp of barium fluoride is 1.0×10–6 at 25 °C. What is the molar solubility of barium fluoride at this temperature?

ANSWER: 6.3×10–3 M Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium 26. What is the Kp expression for the following equilibrium system?

3H2(g) + N2(g) ⇌2NH3(g) a.

b.

c. d.

e.

ANSWER: d 27. At 700 K, hydrogen and iodine exist in equilibrium with hydrogen iodide. What is the value of Kp if the system has a Kc value of 50.5?

H2(g) + I2(g)⇌2HI(g) a. 50.5 b. 25.1 c. 8.23 d. 4.17 e. 75.2 ANSWER: a 28. Which of the following systems would have equal Kc and Kp values? a. 2SO2(g) + O2(g) ⇌2SO3(g) b. 2N2O(g) + O2(g) ⇌4NO(g) c. CO(g) + 2H2(g) ⇌CH3OH(g) d. N2(g)+ 3H2(g) ⇌2NH3(g) e. N2(g) + O2(g) ⇌2NO(g) ANSWER: e Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium 29. At 298 K, the decomposition of dinitrogen tetroxide into nitrogen dioxide reaches equilibrium. What is the value of Kp if the system has a Kc value of 4.63×10–3?

N2O4(g) ⇌2NO2(g) a. 0.0691 b. 0.113 c. 0.280 d. 0.172 e. 0.0752 ANSWER: b 30. The following equilibrium system has a Kp value of 3.0×10–5 at 850 °C. What is the Kc value for the system?

CS2(g) + 4H2(g) ⇌CH4(g) + 2H2S(g) a. 0.91 b. 0.13 c. 0.25 d. 0.19 e. 0.52 ANSWER: c 31. When converting from Kc to Kp, what would be the value of Δn for the following equilibrium system?

C3H8(g) + 5O2(g) ⇌3CO2(g) + 4H2O(g) a. 0 b. 1 c. 2 d. –1 e. –2 ANSWER: b 32. At 298 K, COCl2 reaches equilibrium with CO and Cl2. Given the following Kp value, what is the Kc value for this equilibrium system?

COCl2(g) ⇌CO(g) + Cl2(g)

Kp = 5.5×10–7

a. 5.5×10–7 b. 2.2×10–8 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium c. 5.5×10–8 d. 4.1×10–7 e. 1.8×10–7 ANSWER: b 33. Hydrogen and bromine reach equilibrium with hydrogen bromide. The partial pressures of hydrogen and bromine are both 1.23 atm, and the partial pressure of hydrogen bromide is 3.25×10–4 atm at equilibrium. What is the value of Kp for this system?

H2(g) + Br2(g) ⇌2HBr(g) ANSWER: 6.98×10–8 34. The decomposition of dinitrogen pentoxide into nitrogen dioxide and oxygen reaches equilibrium at 298 K with a Kc equal to 7.5×10−2. What is the value of Kp for this system?

2N2O5(g) ⇌4NO2(g) + O2(g) ANSWER: 1.1×103 35. What is the Kp expression for the following equilibrium system?

NH4SH(s) ⇌NH3(g) + H2S(g) ANSWER: 36. Using the Kc value for reaction (1), what is the Kc value for reaction (2)?

(1) H2(g) + I2(g) ⇌2HI(g) (2) 2HI(g) ⇌H2(g) + I2(g) a. 50.5 b. 101 c. 25.3 d. 0.0198 e. 1.30 ANSWER: d

Kc = 50.5 Kc' = ?

37. Using the Kc value for reaction (1), what is the Kc value for reaction (2)?

(1) N2(g) + 3H2(g) ⇌2NH3(g)

Kc = 6.7×105

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Chapter 14 Equilibrium (2) NH3(g) ⇌N2(g) +

H2(g)

Kc' = ?

a. 3.0×10–6 b. 1.5×10–6 c. 1.2×10–3 d. 3.4×105 e. 2.9×103 ANSWER: c 38. Using the Kc value for reaction (1), what is the Kc value for reaction (2)?

(1) CaCO3(s) ⇌CaO(s) + CO2(g) (2) 3CaCO3(s) ⇌3CaO(s) + 3CO2(g) a. 1.01 b. 0.677 c. 6.09 d. 8.36 e. 1.27 ANSWER: d

Kc = 2.03 Kc' = ?

39. Magnesium carbonate is more soluble in acidic conditions due to the following equilibrium reaction.

MgCO3(s) +H3O+(aq) ⇌Mg2+(aq) + HCO3−(aq) + H2O(l)

Kc = ?

Determine the equilibrium constant for that system using the following information.

MgCO3(s) ⇌Mg2+(aq) + CO32−(aq) HCO3−(aq) + H2O(l) ⇌CO32−(aq) + H3O+(aq) a. 1.5×105

Ksp = 6.82×10−6 Kc = 4.7×10−11

b. 3.2×10−16 c. 6.9×10−6 d. 5.3×104 e. 1.7×109 ANSWER: a 40. Calcium sulfite is more soluble in acidic conditions due to the following equilibrium reaction.

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Chapter 14 Equilibrium CaSO3(s) +H3O+(aq) ⇌Ca2+(aq) + HSO3−(aq) + H2O(l)

Kc = ?

Determine the equilibrium constant for that system using the following information.

CaSO3(s) ⇌Ca2+(aq) + SO32−(aq) HSO3−(aq) + H2O(l) ⇌SO32−(aq) + H3O+(aq) a. 2.5×103

Ksp = 6.8×10−8 Kc = 1.4×10−2

b. 7.2×10−5 c. 4.9×10−6 d. 2.2×104 e. 1.9×10–2 ANSWER: c 41. Barium sulfate is more soluble in acidic conditions due to the following equilibrium reaction.

BaSO4(s) + H3O+(aq) ⇌Ba2+(aq) + HSO4–(aq) + H2O(l)

Kc = ?

Determine the equilibrium constant for that system using the following information.

BaSO4(s) ⇌Ba2+(aq) + SO42−(aq) HSO4−(aq) + H2O(l) ⇌SO42−(aq) + H3O+(aq) ANSWER: 9.2 × 10−9

Ksp = 1.1×10−10 Kc = 1.2×10−2

42. Silver carbonate is more soluble in acidic conditions due to the following equilibrium reaction.

Ag2CO3(s) + H3O+(aq) ⇌2Ag+(aq) + HCO3–(aq) + H2O(l)

Kc = ?

Determine the equilibrium constant for that system using the following information.

Ag2CO3(s) ⇌2Ag+(aq) + CO32–(aq) HCO3–(aq) + H2O(l) ⇌CO32−(aq) + H3O+(aq) ANSWER: 1.9×10–5

Ksp = 8.1×10−12 Kc = 4.3×10−7

43. If a forward reaction has an equilibrium constant Kc, how would the equilibrium constant for the reverse reaction be related to Kc? ANSWER The equilibrium constant for the reverse reaction is the reciprocal (inverse) of : the equilibrium constant for the forward reaction. Mathematically, if Kforward =

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Chapter 14 Equilibrium Kc, then

.

44. Explain how multiplying the coefficients in a chemical reaction by a factor (e.g., 2 or 1/2) affects the equilibrium constant, Kc. ANSWER When you multiply the coefficients of a reaction by a factor, the equilibrium : constant Kc is raised to the power of that factor. For example, if you multiply all the coefficients in a reaction by 2, the new equilibrium constant is Kc' = Kc2.

45. Which of the following statements is correct? a. If Q = K, the reaction will proceed to the right. b. If Q = K, the reaction will proceed to the left. c. If Q < K, the reaction will proceed to the left. d. If Q > K, the reaction will proceed to the right. e. All of the statements are incorrect.

ANSWER: e 46. The concentrations of N2, H2, and NH3 in a mixture are 1.0 M, 3.0 M, and 0.50 M, respectively. What is the reaction quotient (Qc) for this mixture?

N2(g) + 3H2(g) ⇌2NH3(g) a. 0.083 b. 0.25 c. 0.017 d. 0.0093 e. 0.0025 ANSWER: b 47. For a mixture in which [H2] = [F2] =0.40 M and [HF] = 0.20 M, what is the reaction quotient (Qc) for the mixture? Is this mixture at equilibrium? If not, in which direction will the reaction proceed?

H2(g) + F2(g) ⇌2HF(g) Kc = 5.0 a. Qc = 2.0; the reaction proceeds toward products b. Qc = 4.0; the reaction proceeds toward reactants c. Qc = 0.25; the reaction proceeds toward products d. Qc = 8.0; the reaction proceeds toward reactants e. Qc = 5.0; the reaction is at equilibrium ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium 48. For a mixture in which [CO] = 0.50 M, [H2] = 0.20 M, and [CH3OH] = 1.5 M, what is the value of Qc? Is this mixture at equilibrium? If not, in which direction will the reaction proceed?

CO(g) + 2H2(g) ⇌CH3OH(g) Kc = 10 a. Qc = 3.5; the reaction proceeds toward products b. Qc = 15; the reaction proceeds toward reactants c. Qc = 25; the reaction proceeds toward products d. Qc = 75; the reaction proceeds toward reactants e. Qc = 10; the reaction is at equilibrium ANSWER: d 49. In a mixture, the partial pressures of SO2, O2, and SO3 are 0.40 atm, 0.60 atm, and 2.0 atm, respectively.

What is the reaction quotient (Qp)?

2SO2(g) + O2(g) ⇌2 SO3(g) a. 8.3 b. 42 c. 9.1 d. 67 e. 29 ANSWER: b 50. In a mixture, the partial pressure of NO2 is 0.79 atm and the partial pressure of N2O4 is 0.89 atm. What is the value of Qp? Is this mixture at equilibrium? If not, in which direction will the reaction proceed?

2NO2(g) ⇌N2O4(g) Kp = 6.9 a. Qp = 1.4; the reaction proceeds toward products b. Qp = 9.7; the reaction proceeds toward reactants c. Qp = 1.1; the reaction proceeds toward products d. Qp = 27; the reaction proceeds toward reactants e. Qp = 6.9; the reaction is at equilibrium ANSWER: a 51. A reaction mixture initially contains [H2] = 0.10 M, [I2] = 0.10 M, and [HI] = 0.40 M. What is the concentration of HI at equilibrium?

H2(g) + I2(g) ⇌2HI(g)

Kc = 52

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Chapter 14 Equilibrium a. 0.52 M b. 0.36 M c. 0.47 M d. 0.41 M e. 0.75 M ANSWER: c 52. In a mixture, the initial concentrations of H2 and Br2 are each 0.15 M, and the initial concentration of HBr is 0.65 M. What is the equilibrium concentration of HBr? H2(g) + Br2(g) ⇌2HBr(g) a. 0.42 M b. 0.72 M c. 0.87 M d. 0.76 M e. 0.68 M ANSWER: d

Kc = 64

53. Hydrocyanic acid, HCN, ionizes to form H+ and CN– in aqueous solutions. A solution initially contains 0.115 M HCN. What is the concentration of H+ once the system reaches equilibrium?

HCN(aq) ⇌H+(aq) + CN–(aq) a. 8.4×10–6 M

Kc = 6.2×10–10

b. 3.1×10–6 M c. 1.9×10–5 M d. 7.2×10–6 M e. 2.4×10–5 M ANSWER: a 54. If 0.205 mol of both CO and H2O are placed in a 1.0-L reaction vessel, and after the reaction reaches equilibrium there are 0.110 mol of H2, what is the equilibrium constant for the reaction?

CO(g) + H2O(g) ⇌CO2(g) + H2(g) a. 12.1 b. 0.840 c. 0.219 d. 1.34 e. 2.41 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium ANSWER: d 55. What is the equilibrium constant for the following system if the initial concentration of N2 was 5.5×10–3 M, the initial concentration of O2 was 1.05×10–3 M, and the equilibrium concentration of NO is 1.05×10–12 M? N2(g) + O2(g) ⇌2NO(g) ANSWER: 1.9×10–19 56. A mixture of 0.316 M Cl2, 0.459 M F2, and 0.794 M ClF is enclosed in a vessel and heated to 2500 K. Calculate the equilibrium concentration of each gas at 2500 K. Cl2(g) + F2(g) ⇌2ClF(g) Kc = 20.0 at 2500 K ANSWER: [Cl2] = 0.18 M, [F2] = 0.32 M, [ClF] = 1.07 M 57. A mixture contains [CH3COOH] = 0.075 M, [C2H5OH] = 0.085 M, and [CH3COOC2H5] = 0.035 M. What is the reaction quotient for this mixture? Is this mixture at equilibrium? If not, in which direction will the reaction proceed?

CH3COOH(aq) + C2H5OH(aq) ⇌CH3COOC2H5(aq) + H2O(l) ANSWER: Qc = 5.5

Kc = 4.0

Since Qc > Kc, the reaction will proceed in the reverse direction to reach equilibrium.

58. Hydrogen sulfide decomposes to form hydrogen and sulfur. A mixture contains

= 0.150 atm and

=

–5

2.25 × 10 atm. What is the reaction quotient (Qp)? Is this mixture at equilibrium? If not, in which direction will the reaction proceed?

H2S(g) ⇌H2(g) + S(s) ANSWER: Qp = 1.55×10–4

Kp = 1.0×10–7

Since Qp > Kp, the reaction will proceed in the reverse direction to reach equilibrium.

59. Phosphorus pentachloride decomposes to form phosphorus trichloride and chlorine gas. At a certain temperature, the Kc for this reaction is 0.0420. If the initial concentration of PCl5 is 1.00 M and no PCl3 nor Cl2 is initially present, what is the equilibrium concentration of Cl2?

PCl5(g) ⇌PCl3(g) + Cl2(g)

Kc = 0.0420

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Chapter 14 Equilibrium ANSWER: [Cl2] = 0.185 M 60. Formic acid (HCOOH) ionizes in water according to the following equation. The equilibrium constant, Kc, is 1.8×10–4. If the initial concentration of HCOOH is 0.500 M, and no H+ nor HCOO– is initially present, what is the equilibrium concentration of HCOO–?

HCOOH(aq) ⇌H+(aq) + HCOO–(aq) ANSWER: [HCOO–] = 9.4×10–3 M

Kc = 1.8×10–4

61. Which of the following best describes Le Chatelier's principle? a. The equilibrium constant of a reaction is always constant. b. An equilibrium system will shift in the direction that minimizes any disturbances to equilibrium. c. The rate of a reaction is independent of the concentration of reactants. d. The addition of a catalyst will shift the equilibrium towards the products. e. The equilibrium position of a reaction is determined solely by the initial concentrations of the reactants.

ANSWER: b 62. Which of the following actions will produce more PH3?

3KOH(aq) + P4(s) + 3H2O(l) → 3KH2PO2(aq) + PH3(g) a. adding more KH2PO2 to the reaction container b. removing P4 from the reaction container c. removing H2O(l) from the reaction container d. removing KOH from the reaction container e. removing PH3 from the reaction container ANSWER: e 63. In the Haber process for ammonia production, decreasing the pressure will N2(g) + 3H2(g) ⇌2NH3(g) a. shift the equilibrium to the left. b. shift the equilibrium to the right. c. increase the concentration of ammonia. d. increase the value of the equilibrium constant. e. have no effect on the equilibrium.

ANSWER: a Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium 64. What will occur if additional H2 is added to the following equilibrium system?

H2(g) + I2(g) ⇌2HI(g) a. The equilibrium will shift to the left. b. The equilibrium constant will increase. c. The temperature of the system will decrease. d. The concentration of I2 will increase. e. The concentration of HI will increase.

ANSWER: e 65. The evaporation of water in a closed container reaches equilibrium. If the container volume is increased, what is the effect on the equilibrium?

H2O(l) ⇌H2O(g) a. More liquid water will evaporate. b. Some water vapor will condense. c. The equilibrium will not be affected. d. The pressure inside the container will increase. e. The temperature inside the container will decrease.

ANSWER: a 66. Which of the following changes will NOT affect the equilibrium position of a reaction? a. changing the temperature b. changing the pressure for reactions involving gases c. adding a catalyst d. adding additional products e. changing the volume of the container for reactions involving gases ANSWER: c 67. For the reaction CO(g) + 2H2(g) ⇌CH3OH(g), what will occur if the volume of the container is decreased? a. The system will shift to the left. b. The system will shift to the right. c. No shift will occur. d. The equilibrium constant will increase. e. The equilibrium constant will decrease.

ANSWER: b 68. For the reaction CO2(g) + C(s) ⇌2CO(g), what will occur if additional carbon is added? a. The system will shift to the left.

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Chapter 14 Equilibrium b. The system will shift to the right. c. No shift will occur. d. The equilibrium constant will increase. e. The equilibrium constant will decrease.

ANSWER: c 69. What will occur if the temperature is increased for the following equilibrium system?

N2(g) + O2(g) ⇌2NO(g) ΔH > 0 a. The equilibrium will shift to the left. b. The equilibrium will shift to the right. c. No shift will occur. d. The concentration of N2 will increase. e. The concentration of NO will decrease.

ANSWER: b 70. What will occur if the temperature is increased for the following equilibrium system?

CH4(g) + 2O2(g) ⇌CO2(g) + 2H2O(g) ΔH < 0 a. The equilibrium will shift to the left. b. The equilibrium will shift to the right. c. No shift will occur. d. The concentration of CO2 will increase. e. The concentration of O2 will decrease.

ANSWER: a 71. Which of the following will ALWAYS shift the equilibrium to the right? a. increasing the temperature b. decreasing the temperature c. increasing the pressure d. adding a catalyst e. decreasing the concentration of a product ANSWER: e 72. For the reaction 2SO2(g) + O2(g) ⇌2SO3(g), if some SO3 is removed from the system at equilibrium, what will happen to the concentration of O2?

ANSWER: The reaction will shift to the right causing the concentration of O2 to decrease.

73. For the reaction N2(g) + 3H2(g) ⇌2NH3(g), how will decreasing the volume of the container affect the Copyright Macmillan Learning. Powered by Cognero.

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Chapter 14 Equilibrium equilibrium?

ANSWER Decreasing the volume of the container will increase the pressure inside the : container. The equilibrium will shift toward the products since there are fewer moles of gas in the products.

74. The reaction PCl5(g) ⇌PCl3(g) + Cl2(g) is at equilibrium. If the pressure is increased, which way will the equilibrium shift?

ANSWER Increasing the pressure will cause the equilibrium to shift toward the reactants : since there are fewer moles of gas on the reactant side of the equation.

75. How does adding a catalyst affect a system at equilibrium?

ANSWER Adding a catalyst to a system at equilibrium does not change the position of the : equilibrium; it only speeds up the rate at which equilibrium is reached.

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Chapter 13 Kinetics 1. Which of the following would NOT increase the reaction rate of a chemical reaction occurring in an aqueous solution? a. Increasing the temperature. b. Increasing the concentration of a reactant. c. Increasing the surface area of a solid reactant. d. Increasing the amount of the solvent. e. Adding a catalyst.

ANSWER: d 2. Which of the following conditions would likely have the fastest reaction rate for the following chemical reaction?

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) a. [HCl] = 0.050 M at 25 °C b. [HCl] = 0.50 M at 25 °C c. [HCl] = 1.50 M at 50°C d. [HCl] = 0.50 M at 75 °C e. [HCl] = 2.50 M at 75 °C ANSWER: e 3. What is the general reaction rate in terms of the change in concentration of hydrogen?

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) a. b. c. d. e.

ANSWER: b 4. What is the general reaction rate in terms of the change in concentration of INO(g)?

2INO(g) → I2(g) + 2NO(g) a.

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Chapter 13 Kinetics b. c. d. e.

ANSWER: a 5. Aluminum sulfide reacts with water to produce hydrogen sulfide gas and aluminum hydroxide according to the balanced chemical equation shown. If the reaction produces 0.020 mol/L of H2S per second, what is the general reaction rate for this reaction?

Al2S3(s) + 6H2O(l) → 3H2S(aq) + 2Al(OH)3(aq) a. 0.0067 M/s b. 0.010 M/s c. 0.020 M/s d. 0.040 M/s e. 0.060 M/s ANSWER: a 6. What is the general reaction rate for the chemical reaction shown if the concentration of H+ started at 1.5×10– 3 M and was measured to be 5.5×10–4 M after 25.5 s?

IO3−(aq) + 5I−(aq) + 6H+(aq) → 3I2(aq) + 3H2O(l) a. 6.2×10–6 M/s b. 3.7×10–5 M/s c. 4.9×10–6 M/s d. 1.6×10–5 M/s e. 9.1×10–6 M/s ANSWER: a 7. The concentration of H+ started at 2.5×10–3 M and was measured to be 4.2×10–4 M after 35.5 s for the chemical reaction shown. What is the rate of change for I2 over this time period?

IO3−(aq) + 5I−(aq) + 6H+(aq) → 3I2(aq) + 3H2O(l) a. 4.1×10–6 M/s b. 6.2×10–5 M/s Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics c. 9.3×10–6 M/s d. 2.9×10–5 M/s e. 7.0×10–6 M/s ANSWER: d 8. During the reaction H2(g) + I2(g) → 2HI(g), the concentration of H2 was observed to decrease from 0.040 M to 0.010 M in 10 seconds. Calculate the average reaction rate and the rate of formation of HI. ANSWER The average reaction rate is 3.0×10−3 M/s. The rate of formation of HI is : 6.0×10−3 M/s.

9. The concentration of NO2 started at 1.5×10−3 M and decreased to 3.0×10−4 M over a period of 25 seconds for the reaction 2NO2(g)→ 2NO(g) + O2(g). Calculate the average reaction rate and the rate of formation of NO. ANSWER The average reaction rate is 2.4×10−5 M/s. The rate of formation of NO is : 4.8×10−5 M/s.

10. Why does the rate of reaction increase if the temperature is increased?

ANSWER At higher temperatures, molecules move more quickly and collide more : frequently. The collisions also have more energy, making the molecules more likely to react.

11. What is the overall reaction order for a reaction that has the following rate law equation? rate = k[A]2[B] a. 0 b. 1 c. 2 d. 2.5 e. 3 ANSWER: e 12. What are the units for the rate constant for a reaction that has the following rate law equation? rate = k[A]2[B] a. M·s–1 b. M–2·s–1 c. M–1·s–1 d. M–2·s–2 e. s–1 ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics 13. If doubling the concentration of reactant A quadruples the reaction rate, what is the order of the reaction with respect to A? a. zeroth order b. first order c. second order d. third order e. fourth order ANSWER: c 14. For the reaction 2A + B → products, the rate law is determined to be rate = k[A]2[B]. If the concentration of A is tripled and the concentration of B is doubled, by what factor will the rate of the reaction increase? a. 6 b. 12 c. 18 d. 36 e. 72 ANSWER: c 15. An experiment was performed to determine the rate law for the decomposition of X. The concentration of X was varied, and the reaction rate was measured. Plotting the rate as a function of concentration of X yielded the following graph. What is the rate law for this reaction?

a. Rate = k[X]0 b. Rate = k c. Rate = k[X]2 d. Rate = k[X] e. Rate = k[X]3 ANSWER: c 16. An experiment was performed to determine the rate law for the decomposition of X. The concentration of X was varied, and the reaction rate was measured. Plotting the rate as a function of concentration of X yielded the Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics following graph. What is the rate law for this reaction?

a. Rate = k[X]0 b. Rate = k c. Rate = k[X]2 d. Rate = k[X] e. Rate = k[X]3 ANSWER: d 17. Using the experimental data provided, determine the rate law for the following chemical reaction.

2NO2(g) → 2NO(g) + O2(g) Trial Rate (M/s) [NO2] A

0.25 M

1.25×10–2

B

0.35 M

2.45×10–2

C

0.45 M

4.05×10–2

a. Rate = k b. Rate = k[NO2]0 c. Rate = k[NO2]1 d. Rate = k[NO2]2 e. Rate = k[NO2]3 ANSWER: d 18. Using the experimental data provided, determine the rate constant for the following chemical reaction at 25 °C. 2NO2(g) → 2NO(g) + O2(g) Trial Rate (M/s) [NO2] Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics A

15.0 mM

1.55×10–3

B

20.0 mM

2.76×10–3

C

25.0 mM

4.31×10–3

a. 6.89 M–1·s–1 b. 3.10 s–1 c. 0.326 M–1·s–1 d. 4.28 M–1·s–1 e. 0.103 s–1 ANSWER: a 19. An experiment was performed to determine the rate law for the decomposition of X. The concentration of X was varied, and the reaction rate was measured. It was determined that the decomposition of X is a zero-order process. Sketch a plot of the rate as a function of concentration of X that would show the results of this experiment.

ANSWER:

20. Using the experimental data provided, determine the rate law and value of the rate constant for the following chemical reaction.

H2(g) + I2(g) → 2HI(g) Trial

[H2]

[I2]

Rate (M/s)

A

0.100 M

0.100 M

4.65×10–3

B

0.100 M

0.200 M

9.31×10–3

C

0.200 M

0.200 M

1.85×10–2

ANSWER The rate law for the reaction is rate = k[H2][I2]. The value of the rate constant k : is 0.465 M−1·s−1.

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Chapter 13 Kinetics 21. Using the experimental data provided, determine the rate law and value of the rate constant for the following chemical reaction.

NO2(g) + CO(g) → NO(g) + CO2(g) Trial

[NO2]

[CO]

Rate (M/s)

A

0.080 M

0.020 M

5.50×10–3

B

0.080 M

0.040 M

5.52×10–3

C

0.120 M

0.040 M

1.23×10–2

ANSWER The rate law for the reaction is rate = k[NO2]2. The value of the rate constant : is 0.859 M−1·s−1.

22. The decomposition of an organic molecule has a rate constant of 3.25×10–4 s–1 and is a first order reaction.

If the molecule started with a concentration of 0.0375 M, what would the concentration be after 556 s? a. 0.342 M b. 0.215 M c. 0.309 M d. 0.279 M e. 0.313 M ANSWER: e 23. The decomposition of hydrogen peroxide in aqueous solution follows first-order kinetics and has a rate constant of 1.0×10−7 s−1. How long would a 30.0% concentrated solution of hydrogen peroxide take to reach a concentration of 20.0%? a. 42 days b. 47 days c. 13 days d. 79 days e. 28 days ANSWER: b 24. Which of the following is the integrated rate law for a first-order reaction? a. b. c. d. e.

ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics 25. A zeroth-order reaction has a rate constant of 0.025 M/s. If the initial concentration is 0.50 M, how long will it take for the concentration to decrease to 0.10 M? a. 16 s b. 32 s c. 12 s d. 19 s e. 27 s ANSWER: a 26. A second-order reaction has a rate constant of 0.025 mM/s. If the initial concentration of the reactant is 0.55 mM, what will the concentration be after 15 seconds? a. 0.39 mM b. 0.46 mM c. 0.12 mM d. 0.51 mM e. 0.27 mM ANSWER: b 27. A first-order reactant had an initial concentration of 0.600 M. After 45.0 minutes, the concentration of the reactant was 0.300 M. What is the rate constant of the reaction? a. 3.69×10–4 s–1 b. 1.29×10–4 s–1 c. 9.98×10–3 s–1 d. 2.57×10–4 s–1 e. 4.01×10–3 s–1 ANSWER: d 28. The decomposition of cyclohexane follows first-order kinetics and has a rate constant of 1.8×10–2 s–1 at 700 °C. How long would it take 50.0% of a sample of benzene to decompose at that temperature? a. 28 s b. 12 min c. 39 s d. 86 s e. 4.2 min ANSWER: c 29. To determine the order of the reaction A → products, you plot [A], ln[A], and 1/[A] versus time. The plot of 1/[A] versus time produces a straight line. What is the order of the reaction? a. 0 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics b. 1 c. 2 d. 3 e. 4 ANSWER: c 30. The iodine clock reaction follows first-order kinetics and has a rate constant of 2.0×10–2 s–1. If the initial concentration of iodide (I−) is 0.250 M in a 1.0 L container, how many grams of iodine (I2) are produced after 30.0 s?

2I−(aq) + S2O82−(aq) → I2(aq) + 2SO42−(aq) ANSWER: 14 g 31. The decomposition of dinitrogen pentoxide follows second-order kinetics and has a rate constant of 3.38×10–5 s–1. If the initial concentration of dinitrogen pentoxide is 0.125 M in a 500.0 mL container, how many grams of oxygen are produced after 24.0 hours?

2N2O5(g) → 4NO2(g) + O2(g) ANSWER: 0.267 g 32. The decomposition of azomethane follows first-order kinetics and has a rate constant of 3.2×10–5 s–1 at 500 K. How long would it take for half of a sample of azomethane to decompose?

CH3NNCH3(g) → C2H6(g) + N2(g) ANSWER: 22,000 s (6.1 hours) 33. The production of bromine from bromate and bromide in acidic conditions follows first-order kinetics and has a rate constant of 2.00×10–2 s–1. If the initial concentration of bromate (BrO3–) is 0.375 M, what would be its concentration after 30.0 s?

3− − + BrO (aq) + 5Br (aq) + 6H (aq) → 3Br2(aq) + 3H2O(l) ANSWER: 0.206 M 34. What is the half-life of a reactant? a. The time it takes for the reactant concentration to decrease to half of its initial value. b. The time it takes for the reactant concentration to decrease by one-fourth of its initial value. c. The time it takes for the reactant concentration to decrease by 90%. d. The time it takes for the reactant concentration to reach zero. e. The time it takes for the reactant concentration to decrease to one-third of its initial value.

ANSWER: a Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics 35. Iodine-131 is a radioactive isotope that is used to treat hypothyroidism. It has a half-life of approximately eight days. If a hospital received 250 mg of iodine-131, how many milligrams would be present after four weeks? a. 27 mg b. 110 mg c. 22 mg d. 45 mg e. 86 mg ANSWER: c 36. What is the half-life of a first-order reactant with a rate constant equal to 1.25×10–3 s–1? a. 1.98 min b. 2.56 min c. 22.9 s d. 128 s e. 554 s ANSWER: e 37. Radiocarbon dating uses the carbon-14 isotope to determine the age of artifacts containing organic material.

Carbon-14 has a half-life of 5,730 years. What is the rate constant for the decay of the isotope? a. 4.36×10–4 yrs–1 b. 1.21×10–4 yrs–1 c. 5.73×10–4 yrs–1 d. 1.98×10–4 yrs–1 e. 3.47×10–3 yrs–1 ANSWER: b 38. Strontium-90 has a half-life of 28.8 years. How many years would it take for a sample to decay to 12.5% of its original amount? a. 173 years b. 144 years c. 115 years d. 86.4 years e. 57.6 years ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics 39. A first-order reaction has a half-life of 5.00 minutes. What will be the concentration of the reactant after 20.0 minutes if the initial concentration is 1.00 M? a. 0.500 M b. 0.250 M c. 0.125 M d. 0.0625 M e. 0.0313 M ANSWER: d 40. For a first-order reaction, the half-life is a. dependent on the initial concentration of the reactant. b. constant and independent of the concentration of the reactant. c. longer for faster reactions. d. the same as the rate constant. e. shorter for slower reactions.

ANSWER: b 41. If the rate constant of a first-order reaction is 2.00×10−3 s−1, what is the half-life of the reaction?

ANSWER: 347 s 42. A certain radioactive isotope has a half-life of 5.00 minutes. What percentage of the isotope will remain after 15.0 minutes?

ANSWER: 12.5% 43. Carbon-14 has a half-life of 5,730 years. How old is a wooden artifact if it contains 6.25% of its original carbon-14?

ANSWER: 22,900 years 44. Which of the following best describes a reaction mechanism? a. The overall equation that shows the reactants and products of a reaction. b. A detailed sequence of elementary steps that describe how reactants are converted into products. c. The energy change that occurs during a chemical reaction. d. The rate at which reactants are consumed and products are formed. e. The temperature at which a reaction occurs.

ANSWER: b 45. What is an elementary step in a reaction mechanism? a. A single event or step in a reaction that shows a change in bonding.

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Chapter 13 Kinetics b. The overall chemical reaction equation. c. The fastest step in a reaction mechanism. d. The slowest step in a reaction mechanism. e. The final product formation step.

ANSWER: a 46. A unimolecular elementary step involves a. a catalyst accelerating the reaction. b. the formation of an intermediate complex. c. three molecules interacting simultaneously. d. two molecules colliding to form products. e. one molecule decomposing or rearranging.

ANSWER: e 47. Which of the following best describes a bimolecular elementary step? a. The reaction of a single molecule. b. The interaction between a molecule and a catalyst. c. The collision and reaction between two molecules. d. The simultaneous collision of three molecules. e. The rearrangement of atoms within a molecule.

ANSWER: c 48. What is the rate-determining step in a reaction mechanism? a. The fastest step in the reaction sequence. b. The step with the lowest activation energy. c. The first step in the reaction mechanism. d. The slowest step in the reaction sequence. e. The step that involves the formation of the final product.

ANSWER: d 49. What is the rate law for the following reaction, given the elementary steps below?

Overall Reaction: CH4 + Cl2 → CH3Cl + HCl Mechanism: Cl2 → 2Cl (slow) CH4 + Cl → CH3 + HCl (fast) CH3 + Cl → CH3Cl (fast) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics a. Rate = k[CH4][Cl2] b. Rate = k[CH4] c. Rate = k[Cl2] d. Rate = k[CH4][Cl2]2 e. Rate = k[CH3][Cl]2 ANSWER: c 50. What is the rate law for the following reaction, given the elementary steps below?

Overall Reaction: 2NO2 + F2 → 2NO2F Mechanism: NO2 + F2 → NO2F + F (slow) NO2 + F → NO2F (fast) a. Rate = k[NO2][F2] b. Rate = k[NO2]2[F2] c. Rate = k[NO2][F2]2 d. Rate = k[NO2] e. Rate = k[NO2] ANSWER: a 51. What is the molecularity of the following elementary step of a chemical reaction?

NO2 + NO2 → NO3 + NO ANSWER: It is a bimolecular elementary step.

52. What is the molecularity of the following elementary step of a chemical reaction?

F2 → 2F ANSWER: It is a unimolecular elementary step.

53. What is the rate law for the following reaction, given the elementary steps below?

Overall Reaction: NO2 + CO → NO + CO2 Mechanism: Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics NO2 + NO2 → NO3 + NO (slow) NO3 + CO → NO2 + CO2 (fast) ANSWER: rate = k[NO2]2 54. What is the rate law for the following reaction, given the elementary steps below?

Overall Reaction: H2 + 2ICl → 2HCl + I2 Mechanism: H2 + ICl → HI + HCl (slow) HI + ICl → HCl + I2 (fast) ANSWER: rate = k[H2][ICl] 55. On a reaction energy diagram, what does the difference in energy between the reactants and the peak of the diagram represent? a. The equilibrium constant of the reaction. b. The activation energy of the reaction. c. The change in energy of the reaction. d. The slowest step in the reaction sequence. e. The rate constant of the reaction.

ANSWER: b 56. What information CANNOT be obtained from a reaction energy diagram? a. The number of elementary steps for the reaction. b. The activation energy of the reaction. c. The change in energy of the reaction. d. The relative energies of reactants, intermediates, and products e. The rate constant of the reaction.

ANSWER: e 57. Which of the following best describes the transition state in a chemical reaction? a. The stable intermediate formed during the reaction. b. The highest energy state along the reaction path, where bonds are partially formed and broken. c. The final product of the reaction. d. The initial reactants before the reaction begins. e. The step in the reaction mechanism with the lowest energy.

ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics 58. A reaction has an activation energy of +60.0 kJ/mol and a net energy change of –45.2 kJ/mol. Does this reaction absorb or release energy? Will the reaction occur at room temperature a. The reaction releases energy and will occur at room temperature. b.The reaction absorbs energy and will occur at room temperature. c. The reaction absorbs energy and it will not occur at room temperature. d.The reaction releases energy and it will not occur at room temperature. e. The reaction releases energy, but it is not possible to determine whether or not it will occur at room temperature.

ANSWER: d 59. Activation energy is defined as a. the total energy released during a reaction. b. the energy difference between reactants and products. c. the energy barrier for a reaction. d. the energy required to break all bonds in the reactants. e. the energy absorbed by the reactants during the reaction.

ANSWER: c 60. In the Arrhenius equation, the frequency factor (A) represents a. the temperature dependence of the reaction rate. b. the fraction of molecules with enough energy to react. c. the total number of reactant molecules. d. the probability that the particles will collide in the correct orientation. e. the change in energy of the reaction.

ANSWER: d 61. The reaction energy diagram for a reaction is shown. What is the approximate activation energy for this reaction?

a. 70 kJ/mol Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics b. 40 kJ/mol c. 20 kJ/mol d. 80 kJ/mol e. 50 kJ/mol ANSWER: e 62. The rate constant of a chemical reaction was determined at several temperatures. A graph of ln(k) versus 1/T (K–1) produced a straight line. The equation of the best-fit line is y = –75200x + 15.2. What are the activation energy and the frequency factor of this reaction?

ANSWER The activation energy of the reaction is 625 kJ/mol, and the frequency factor is : 3.99×106 M–1·s–1.

63. The rate constant of a chemical reaction was determined at several temperatures. A graph of ln(k) versus 1/T (K–1) produced a straight line. The equation of the best-fit line is y = –11200x + 19.8. What is the activation energy of this reaction?

ANSWER: The activation energy of the reaction is 93.1 kJ/mol.

64. The reaction energy diagram for a reaction is shown. What is the approximate activation energy for this reaction?

ANSWER: The activation energy of the reaction is approximately 20 kJ/mol.

65. The reaction energy diagram for a reaction is shown. What are the approximate activation energy and net energy change for this reaction?

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Chapter 13 Kinetics

ANSWER The activation energy of the reaction is approximately 30 kJ/mol, and net energy change is approximately –30 kJ/mol. : 66. A gas phase reaction has an activation energy of 50.0 kJ/mol. At 30.0 °C, the reaction takes place at a rate of 0.200 M−1 min−1. What is the rate of this reaction if the temperature is decreased to 10.0 °C? ANSWER: 0.0492 M−1 min−1 67. How do catalysts change the rate of a reaction? a. By increasing the temperature of the reaction. b. By providing a lower activation energy pathway. c. By increasing the concentration of reactants. d. By increasing the pressure of the system. e. By altering the stoichiometry of the reaction.

ANSWER: b 68. Which statement is TRUE about homogeneous catalysts? a. They exist in the same phase as the reactants. b. They are always in the solid phase. c. They exist in a different phase than the reactants. d. They are a reaction intermediate. e. They increase the activation energy of the reaction.

ANSWER: a 69. Which of the following acts as a biological catalyst to speed up biochemical reactions in living organisms? a. DNA b. hormones c. lipids d. steroids e. enzymes Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics ANSWER: e 70. Which statement is TRUE about heterogeneous catalysts? a. They exist in the same phase as the reactants. b. They are always in the solid phase. c. They exist in a different phase than the reactants. d. They are a reaction intermediate. e. They increase the activation energy of the reaction.

ANSWER: c 71. Partially hydrogenated oils are produced by the hydrogenation of vegetable oils. This process uses a metal catalyst, typically nickel, to add hydrogen to the carbon-carbon double bonds in the oils, converting them to single bonds and resulting in a semi-solid form of fat. What type of catalysis does this represent? a. homogeneous catalysis b. heterogeneous catalysis c. enzyme catalysis d. photon-induced catalysis e. acid-base catalysis ANSWER: b 72. Sucrase catalyzes the hydrolysis of sucrose into glucose and fructose in the digestive system. Sucrase binds to sucrose and facilitates its breakdown through a specific active site interaction. What type of catalysis does this represent? a. homogeneous catalysis b. heterogeneous catalysis c. enzyme catalysis d. photon-induced catalysis e. acid-base catalysis ANSWER: c 73. What is the overall reaction and catalyst for the following reaction mechanism?

NO + O2 → NO2

NO2 + SO2 → NO + SO3 ANSWER: The overall reaction is SO2 +

O2 → SO3

The catalyst is NO. Copyright Macmillan Learning. Powered by Cognero.

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Chapter 13 Kinetics 74. What is the overall reaction and catalyst for the following reaction mechanism?

NO2 + NO2⟶ NO3 + NO NO3 + CO⟶ CO2 + NO2 ANSWER: The overall reaction is NO2 + CO ⟶ CO2 + NO The catalyst is NO2.

75. Sketch an energy diagram for an uncatalyzed exothermic reaction and the same reaction with a catalyst that alters the reaction mechanism. In your sketch, label both axes, the two pathways, the activation energies for both reactions, and the net energy change.

ANSWER:

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Chapter 12 Solutions 1. A 750.0 mL bottle of wine contains 12.9% (v/v) alcohol. How many milliliters of alcohol are in the bottle? a. 24.5 mL b. 32.6 mL c. 96.8 mL d. 15.9 mL e. 4.23 mL ANSWER: c 2. What volume of ethanol does a 20.0-ounce bottle of hand sanitizer contain with an ethanol concentration of 62.5% (v/v)? a. 13.6 oz b. 11.8 oz c. 7.50 oz d. 17.8 oz e. 12.5 oz ANSWER: e 3. How many moles of HCl are present in a 5.00 L bottle of 37.0% (m/m) hydrochloric acid? The density of the solution is 1.19 g/mL. a. 60.4 mol b. 20.5 mol c. 17.8 mol d. 8.98 mol e. 35.5 mol ANSWER: a 4. What is the concentration (in ppb) of a 150.0 mL solution that contains 2.65 μg of lead? a. 0.0177 ppb b. 17.7 ppb c. 12.6 ppb d. 0.0126 ppb Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions e. 187 ppb ANSWER: b 5. What is the concentration (in ppm) of a 250.0-mL solution that contains 4.56×10–3 g of mercury? Assume the density of the solution is the same as water. a. 18.2 ppm b. 227 ppm c. 17.5 ppm d. 126 ppm e. 87.9 ppm ANSWER: a 6. What is the name of the minor component in a homogenous mixture? a. solvent b. solution c. solute d. solid phase e. liquid phase ANSWER: c 7. What is the name of the major component of a homogenous mixture? a. solvent b. solution c. solute d. solid phase e. liquid phase ANSWER: a 8. Which term refers to the amount of solute present in a solution? a. solvent b. solution c. solute d. concentration e. dilute ANSWER: d 9. What is the formula for calculating the molality of a solution? a.

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Chapter 12 Solutions b. c. d. e.

ANSWER: e 10. What is the molality of a solution prepared by dissolving 45.6 g CaCl2 in 750.0 g of water? a. 0.275 m b. 0.548 m c. 0.0608 m d. 0.113 m e. 0.783 m ANSWER: b 11. A 1.0-gallon bottle of white distilled vinegar has an acetic acid concentration of 5.0% (m/m). How many grams of acetic acid are present in the bottle? Assume the density of vinegar is 1.00 g/mL. (1 gal = 3.785 L) a. 160 g b. 210 g c. 85 g d. 190 g e. 50 g ANSWER: d 12. What is the concentration of a 22.0% (m/m) HCl solution in molarity? The density of the solution is 1.11 g/mL. a. 4.80 M b. 12.2 M c. 1.75 M d. 6.70 M e. 3.68 M ANSWER: d 13. How many grams of KNO3 must be dissolved in 250.0 g of water to make a 0.175 m solution? a. 5.87 g b. 4.42 g Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions c. 3.20 g d. 7.34 g e. 5.92 g ANSWER: b 14. A hard water sample has a calcium carbonate concentration of 565 ppm. How many milligrams of calcium carbonate are present in a 250.0 mL hard water sample? Assume the solution has the same density as water.

ANSWER: 141 mg 15. How many micrograms of lead are found in a 125 mL water sample tested to have a 12 ppb lead concentration? Assume the solution has the same density as water.

ANSWER: 1.5 μg 16. What is the molarity of a 5.0% (m/m) HCl solution? The density of the solution is 1.02 g/mL. ANSWER: 1.4 M 17. What is the molality of a solution prepared by adding 45.5 g of ethylene glycol (C2H6O2) to 750.0 mL of water? The density of water is 1.00 g/mL. ANSWER: 0.977 m 18. What is the mole fraction of methanol (CH3OH) in a solution prepared by adding 95.6 g of methanol to 465 mL of ethanol (CH3CH2OH)? The density of ethanol is 0.789 g/mL. ANSWER: 0.272 19. How much energy is required to dissolve 25.6 g of KCl in 350 mL of water?

KCl(s) → KCl(aq) = +17.7 kJ a. 6.08 kJ b. 7.23 kJ c. 12.6 kJ d. 8.82 kJ e. 4.71 kJ ANSWER: a 20. What factor causes compounds with a slightly endothermic dissolution process to dissolve spontaneously? a. change in pressure b. change in temperature c. change in enthalpy d. change in entropy e. Endothermic dissolution processes do not dissolve spontaneously.

ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions 21. Which of the following processes would increase the solubility of a gas in a solution the most? a. increasing the pressure and decreasing the temperature b. decreasing the pressure and decreasing the temperature c. increasing the pressure and increasing the temperature d. decreasing the pressure and increasing the temperature e. increasing the temperature at constant pressure ANSWER: a 22. Which of the following gases would have the highest solubility in water? a. CH4 b. HF c. F2 d. CO e. H2 ANSWER: b 23. Which of the following compounds is a nonelectrolyte when dissolved in water? a. NaNO3 b. HNO3 c. NO2 d. NH4Cl e. HCl ANSWER: c 24. What is the solubility of oxygen in water at a pressure of 1.14 atm? Henry’s law constant for oxygen is 1.2×10–3 M/atm. a. 2.6×10–3 M b. 1.4×10–3 M c. 1.7×10–3 M d. 6.2×10–3 M e. 9.8×10–4 M ANSWER: b 25. What pressure is required to dissolve 14.3 g of CO2 in 2.5 L of water at 25 °C? Henry’s law constant for carbon dioxide is 3.3×10–2 M/atm. a. 2.9 atm Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions b. 4.2 atm c. 1.8 atm d. 3.9 atm e. 3.1 atm ANSWER: d 26. What is the maximum amount of oxygen that can dissolve in a 750.0 mL sample of water at 1.02 atm and 25 °C? Henry’s law constant for oxygen is 1.2×10–3 M/atm. a. 39.2 mg b. 58.8 mg c. 29.4 mg d. 14.7 mg e. 19.6 mg ANSWER: c 27. Which of the following gases is most soluble in water at a pressure of 1 atm and a temperature of 25 °C? a. oxygen b. nitrogen c. ozone d. carbon dioxide e. chlorine ANSWER: e 28. Describe the three steps, in terms of energy, that occur when an ionic compound dissolves in water.

ANSWER Energy is needed to overcome the ionic bonds within the crystal lattice of the : solid compound, separating the ions. Energy is also required to overcome the hydrogen bonding between water molecules to create space for the dissolved ions. Energy is released when the separated ions interact with water molecules, forming ion-dipole interactions. This step releases energy as the ions are stabilized by being surrounded by water molecules.

29. Describe how pressure and temperature affect the solubility of gases in an aqueous solution.

ANSWER Increasing pressure increases the solubility of gases in aqueous solutions while : increasing temperature generally decreases gas solubility in an aqueous solution.

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Chapter 12 Solutions 30. Why does increasing pressure increase the solubility of gases in a solution?

ANSWER When the pressure of a gas above the solution is increased, more gas particles : are forced into the solution, increasing solubility. Conversely, gas particles can escape if the pressure is decreased, and the solubility decreases.

31. Why does the solubility of gases in a solution decrease as the temperature of the solution increases?

ANSWER As temperature rises, the kinetic energy of gas particles increases, making them : more likely to escape from the solution into the gas phase. This results in a lower solubility of the gas in the liquid.

32. The solubility of argon in water is 62 mg/L at 20 °C and 760 torr. What is the Henry’s law constant for argon?

ANSWER: 0.0016 M/atm 33. What is the total ion concentration of a 0.175 M Ca(NO3)2 solution? Assume the solution is an ideal electrolyte solution. a. 0.0583 M b. 0.175 M c. 0.350 M d. 0.525 M e. 0.700 M ANSWER: d 34. Which of the following solutions has the highest nitrate concentration? Assume each solution is an ideal electrolyte solution. a. [Ca(NO3)2] = 0.050 M b. [AgNO3] = 0.100 M c. [Mg(NO3)2] = 0.080 M d. [Au(NO3)3] = 0.060 M e. [KNO3] = 0.120 M ANSWER: d 35. Which of the following solutions have the highest chloride concentration? Assume each solution is an ideal electrolyte solution. a. [CaCl2] = 0.120 M b. [NaCl] = 0.070 M c. [MgCl2] = 0.090 M Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions d. [FeCl3] = 0.050 M e. [KCl] = 0.170 M ANSWER: a 36. What factor reduces the number of particles that are dissolved in a solution? a. dissociation b. ion-dipole interactions c. hydrogen bonding d. solubility e. ion pairs ANSWER: e 37. A nitric acid (HNO3) solution has a concentration of 0.500 M. What is the total ion concentration of the solution given a van't Hoff factor of 1.9? a. 0.50 M b. 1.0 M c. 0.95 M d. 1.9 M e. 1.5 M ANSWER: c 38. A potassium carbonate (K2CO3) solution has a concentration of 1.50 M. What is the total ion concentration of the solution given a van't Hoff factor of 2.5? a. 4.5 M b. 3.0 M c. 2.6 M d. 3.8 M e. 1.5 M ANSWER: d 39. A magnesium chloride (MgCl2) solution has a concentration of 0.50 M. What is the difference in the total ion concentration between the ideal and real solutions given a van't Hoff factor of 2.6? a. 0.20 M b. 0.30 M c. 0.40 M Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions d. 0.10 M e. 0.25 M ANSWER: a 40. A potassium sulfate (K2SO4) solution has a concentration of 0.175 M. What is the difference in the total ion concentration between the ideal and real solutions given a van't Hoff factor of 2.4? a. 0.42 M b. 0.11 M c. 0.21 M d. 0.075 M e. 0.24 M ANSWER: b 41. Which property of real solutions causes a lower ion concentration compared to the concentration of an ideal electrolyte solution?

ANSWER In real solutions, ions of opposite charges can attract each other and form ion : pairs. This reduces the number of dissociated ions in the solution compared to what is expected in an ideal electrolyte solution where complete dissociation is assumed.

42. What is the significance of the van’t Hoff factor?

ANSWER The van’t Hoff factor describes the number of particles that form per formula : of the solute unit in a real solution.

43. A potassium hydroxide solution has a concentration of 2.50 M. What is the difference in total ion concentration between the ideal and real solutions given a van't Hoff factor of 1.8?

ANSWER: 0.50 M 44. A hydrochloric acid solution has a concentration of 6.00 M. What is the total ion concentration of the solution given a van't Hoff factor of 1.9?

ANSWER: 11 M 45. Which of the following is a colligative property of solutions? a. density b. surface tension c. boiling point elevation d. viscosity e. specific gravity ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions 46. Which statement best describes colligative properties? a. They depend on the chemical identity of the solute particles. b. They depend on the number of solute particles in the solution. c. They depend only on the temperature of the solution. d. They are independent of the concentration of the solution. e. They depend on the charge of the solute particles.

ANSWER: b 47. What is the effect on the freezing point of water when a nonvolatile solute is added? a. It increases. b. It decreases. c. It remains unchanged. d. It depends on the molar mass of the solute. e. It depends on the charge of the solute particles.

ANSWER: b 48. Which of the following will result in the greatest boiling point elevation when dissolved in water? a. 1 mol NaCl b. 1 mol C6H12O6 (glucose) c. 1 mol MgCl2 d. 1 mol HC2H3O2 (acetic acid) e. 1 mol CO2 ANSWER: c 49. Which solution will have the lowest freezing point? a. 1.0 m NaCl b. 1.0 m C6H12O6 (glucose) c. 1.0 m Ca(NO3)2 d. 1.0 m HCl e. 1.0 m NH3 ANSWER: c 50. The osmotic pressure of a solution is 5.0 atm at 25 °C. What will the osmotic pressure be if the concentration of the solution is doubled? a. 2.5 atm b. 5.0 atm c. 7.5 atm d. 10.0 atm Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions e. 20.0 atm ANSWER: d 51. What is the osmotic pressure of a 0.500 M KCl solution at a temperature of 25.0 °C? For KCl, i = 1.8. a. 22 atm b. 11 atm c. 1.8 atm d. 7.5 atm e. 13 atm ANSWER: a 52. What is the concentration of a KNO3 solution with an osmotic pressure of 10.3 atm at a temperature of 25.0 °C? For KNO3, i = 1.8. a. 0.42 M b. 2.8 M c. 5.0 M d. 4.8 M e. 0.24 M ANSWER: e 53. Which of the following solutions will have the highest osmotic pressure at room temperature? a. 4.0 M NaCl b. 1.5 M CaCl2 c. 2.0 M Mg(NO3)2 d. 2.5 M HNO3 e. 3.5 M HCl ANSWER: a 54. The tendency of water to move toward regions of higher solute concentration is caused by which of the following properties? a. osmotic pressure b. enthalpy c. surface tension d. vapor pressure e. specific gravity ANSWER: a 55. A solution is prepared by dissolving 4.00 g of an unknown nonelectrolyte compound in 100.0 g of water.

The freezing point of the solution is found to be –1.24 °C. Given that the freezing point depression constant (Kf) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions for water is 1.86 °C/m, calculate the molar mass of the unknown compound. a. 22.4 g/mol b. 24.8 g/mol c. 49.6 g/mol d. 59.9 g/mol e. 99.2 g/mol ANSWER: d 56. What is the boiling point elevation of a solution prepared by adding 55.0 g of KCl to 750.0 g of ethanol?

Use the following information for the solution: Kb = 1.23 °C/m and i = 1.8.

ANSWER: 2.2 °C 57. What is the freezing point of a solution prepared by adding 25.0 g of NaCl to 500.0 g of water? Use the following information for the solution: Kf = 1.86 °C/m and i = 1.8.

ANSWER: –2.87 °C 58. How many grams of sodium chloride are needed to reduce the freezing point of 10.0 L of water by 5.0 C? For water, d = 1.00 g/mL and Kf 1.86 °C/m. For sodium chloride, i = 1.8.

ANSWER: 874 g 59. A solution is made by dissolving 10.0 g of an unknown nonelectrolyte compound in 100.0 g of benzene. The solution's freezing point is 3.77 °C lower than that of pure benzene. Given that the freezing point depression constant (Kf) for benzene is 5.12 °C/m, calculate the molar mass of the unknown compound.

ANSWER: 136 g/mol 60. A solution of 3.50 g of an unknown nonelectrolyte solute in 200.0 g of chloroform (CHCl3) raises the boiling point of chloroform by 1.24 °C. Given that the boiling point elevation constant (Kb) for chloroform is 3.63 °C/m, calculate the molar mass of the unknown solute.

ANSWER: 51.3 g/mol 61. Which of the following statements is true regarding the vapor pressure of a solution containing a nonvolatile solute? a. It is higher than the vapor pressure of the pure solvent. b. It is lower than the vapor pressure of the pure solvent. c. It is equal to the vapor pressure of the pure solvent. d. It depends on the charge of the solute particles. e. It depends on the molar mass of the solute particles.

ANSWER: b 62. Which of the following interactions leads to deviations from Raoult’s law, where an increase in the vapor Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions pressure of a solution containing a volatile substances is observed? a. strong solvent-solute interactions b. weak solvent-solute interactions c. strong solute-solute interactions d. weak solute-solute interactions e. strong solvent-solvent interactions ANSWER: b 63. Which of the following interactions leads to deviations from Raoult’s law, where a decrease in the vapor pressure of a solution containing a volatile substance is observed? a. strong solvent-solute interactions b. weak solvent-solute interactions c. strong solute-solute interactions d. weak solute-solute interactions e. strong solvent-solvent interactions ANSWER: a 64. The vapor pressure of water at 25 °C is 23.8 torr. What is the vapor pressure of a solution prepared by adding 4.50 moles of nonvolatile nonelectrolyte compound to 750.0 mL of water? Assume the density of water is 1.00 g/mL. a. 19.3 torr b. 17.8 torr c. 22.6 torr d. 21.5 torr e. 20.1 torr ANSWER: d 65. The solubility of NaCl in water is 360 g/L at 25 °C? What is the vapor pressure of a saturated solution of NaCl? Assume the density of water is 1.00 g/mL. The vapor pressure of water at 25 °C is 23.8 torr. a. 19.5 torr b. 16.8 torr c. 18.7 torr d. 20.3 torr e. 21.1 torr ANSWER: d 66. The vapor pressure of a solution is 186 torr at 35 °C. What is the vapor pressure of the pure solvent if the mole fraction of the solvent is 0.80? a. 150 torr b. 190 torr c. 230 torr Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions d. 250 torr e. 260 torr ANSWER: c 67. A chemist prepares a solution containing a nonelectrolyte dissolved in acetone. What is the vapor pressure of the solution if the mole fraction of acetone is 0.750? The vapor pressure of acetone is 421 torr at 40.0 °C. a. 561 torr b. 316 torr c. 389 torr d. 375 torr e. 298 torr ANSWER: b 68. The vapor pressure of acetone is 230.0 torr, and the vapor pressure of water is 23.8 torr at 25 °C. What is the vapor pressure of 50/50 mole fraction acetone/water solution? a. 206 torr b. 254 torr c. 178 torr d. 127 torr e. 153 torr ANSWER: d 69. A hand sanitizer solution contains 60.0% (m/m) ethanol (CH3CH2OH) in water. What is the vapor pressure of the solution? The vapor pressure of ethanol is 58.7 torr and the vapor pressure of water is 23.8 torr at 25 °C? a. 45.2 torr b. 36.7 torr c. 56.9 torr d. 113 torr e. 72.9 torr ANSWER: b 70. Which of the following methods is best suited for separating a mixture of two liquids with different boiling points? a. filtration b. chromatography c. distillation d. decantation e. centrifugation ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 12 Solutions 71. The vapor pressure of acetone is 230 torr, and the vapor pressure of water is 23.8 torr at 25 °C. What mole fraction of acetone would result in a solution with a vapor pressure of 180 torr?

ANSWER: 0.76 72. The vapor pressure of acetone is 230 torr, and the vapor pressure of water is 23.8 torr at 25 ºC. What mole fraction of acetone would result in a solution with a vapor pressure of 180 torr?

ANSWER: 0.76 73. An aqueous solution containing a 0.15-mole fraction of a nonvolatile, nonelectrolyte compound has a vapor pressure of 58.9 torr at a specific temperature. What is the vapor pressure of water at that temperature?

ANSWER: 69 torr 74. What causes the vapor pressure of a solution containing a nonvolatile solute to be lower than the vapor pressure of the pure solvent?

ANSWER When a nonvolatile solute is dissolved in a solvent, the solute particles occupy : space at the liquid's surface. This reduces the number of solvent molecules at the surface that can escape into the vapor phase, leading to a decrease in the solution's vapor pressure 75. How would the vapor pressure of a complex mixture of volatile substances, such as crude oil, be determined?

ANSWER The vapor pressure of the solution would be the sum of the partial vapor : pressures of each component in the solution.

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Chapter 11 Intermolecular Forces, Liquids, and Solids 1. Which of the following intermolecular forces is the weakest? a. ion-dipole interactions b. hydrogen bonding c. dipole-dipole interactions d. ionic bonds e. London dispersion forces ANSWER: e 2. Which type of intermolecular force occurs between all molecules, regardless of their polarity? a. hydrogen bonding b. ion-dipole interactions c. London dispersion forces d. ionic bonds e. dipole-dipole interactions ANSWER: c 3. What is the strongest intermolecular force present in a sample of H2O at room temperature? a. hydrogen bonding b. metallic bonding c. London dispersion forces d. ionic bonding e. dipole-dipole interactions ANSWER: a 4. Which of the following compounds will exhibit hydrogen bonding in the liquid state? a. CH4 b. HBr c. NH3 d. CO2 e. CH2O ANSWER: c 5. Which of the following molecules has the highest boiling point? a. CO2 b. H2S c. CH2F2 d. C3H8 e. CH3CH2OH Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids ANSWER: e 6. Which intermolecular force is primarily responsible for the high boiling point of water compared to other molecules of similar size? a. hydrogen bonding b. ion-dipole interactions c. London dispersion forces d. covalent bonding e. dipole-dipole interactions ANSWER: a 7. Which of the following statements is TRUE regarding London dispersion forces? a. They are only present in polar molecules. b. They are the result of instantaneous dipoles. c. They are stronger than hydrogen bonds. d. They do not occur in noble gases. e. They are a type of ionic bond.

ANSWER: b 8. Which type of intermolecular force is primarily responsible for the interactions between molecules of I2 in the solid state? a. hydrogen bonding b. ion-dipole interactions c. London dispersion forces d. covalent bonding e. dipole-dipole interactions ANSWER: c 9. Which of the following bonds would most likely exhibit hydrogen bonding in a molecule? a. H–C b. H–P c. H–N d. H–S e. H–Br ANSWER: c 10. What intermolecular forces are present in a sample of acetic acid?

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Chapter 11 Intermolecular Forces, Liquids, and Solids

ANSWER Acetic acid exhibits London dispersion forces, dipole-dipole interactions, and : hydrogen bonding.

11. Rank the following molecules in order of increasing strength of intermolecular forces: CH2O, CO2, CH4, CH3OH. ANSWER: The order of increasing strength of intermolecular forces is: CH4 < CO2 < CH2O < CH3OH

12. Why is hydrogen bonding a much stronger intermolecular force than dipole-dipole interactions?

ANSWER When hydrogen covalently bonds to a much more electronegative element (F, : O, or N), its only electron is pulled away, exposing the positive nucleus, which results in a higher partial positive charge. All other elements have more electrons that can shield the nucleus; thus, they do not develop as high of a partial positive charge when covalently bonded to a much more electronegative element. The larger partial positive charge increases the attraction of the hydrogen atom to neighboring molecules with slightly negative atoms. This stronger attraction brings those molecules close enough to form an intermolecular force that is almost the same distance as a covalent bond.

13. Which covalent bonds lead to hydrogen bonding between molecules?

ANSWER: Hydrogen bonds occur between molecules containing H–F, H–O, or H–N bonds.

14. Explain how London dispersion forces occur between two nonpolar molecules.

ANSWER London dispersion forces arise from two processes: instantaneous dipoles and : induced dipoles. Instantaneous dipoles occur when the electron distribution within a nonpolar molecule becomes momentarily asymmetrical, creating a temporary partial charge. This temporary dipole can induce a similar dipole in a neighboring molecule, resulting in a weak, transient attraction between the molecules.

15. When mixed, which of the following combinations would exhibit ion-dipole forces? a. CH3OH(l) and H2O(l) b. C12H22O11(s) and H2O(l) c. NH3(g) and O2(g) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids d. NaOH(s) and H2O(l) e. HF(g) and NH3(g) ANSWER: d 16. In which of the following pairs of molecules would dipole-dipole interactions be the most significant? a. CH3Cl and CH4 b. HCl and HF c. Cl2 and Br2 d. CO2 and CH4 e. O2 and N2 ANSWER: b 17. Which of the following compounds would have the lowest solubility in ethanol, CH2CH3OH? a. CH3OH b. C12H22O11 c. CH2NH3 d. H2O e. C6H6 ANSWER: e 18. Which of the following compounds would have the highest solubility in benzene, C6H6? a. CH3CH2OH b. C8H18 c. CH2NH3 d. H2O e. CH3Cl ANSWER: b 19. Which of the following solvents would dissolve a sample of nonpolar wax compounds? a. CH3CH2OH b. HCON(CH3)2 c. CH3CN d. H2O e. CCl4 ANSWER: e Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids 20. Which intermolecular force is primarily responsible for dissolving ionic compounds in water? a. hydrogen bonding b. ion-dipole forces c. London dispersion forces d. covalent bonding e. dipole-dipole interactions ANSWER: b 21. Which of the following solvents would be the most effective for dissolving a nonpolar compound? a. CH3CH2CH2CH2CH3 b. CH3COCH3 c. CH3Cl d. CH3SOCH3 e. CH3CN ANSWER: a 22. Describe the structure of a molecule that can be used as a detergent.

ANSWER A molecule with a long nonpolar hydrocarbon chain and a polar head group : can be used as a detergent.

23. What does the phrase “like dissolves like” imply regarding the solubility of a compound in a solvent?

ANSWER The phrase "like dissolves like" implies that a solute will dissolve best in a : solvent with a similar polarity. Polar solutes tend to dissolve well in polar solvents, while nonpolar solutes dissolve well in nonpolar solvents.

24. What is a micelle?

ANSWER Micelles are a cluster of molecules that form a spherical assembly with the : polar heads oriented on the outside and the nonpolar tails orientated to the cluster's center.

25. Describe the ion-dipole interactions that enable an ionic compound like NaCl to dissolve in water.

ANSWER The positive sodium ions (Na⁺) and negative chloride ions (Cl⁻) interact with : the polar water molecules. The partially negative oxygen atoms in water molecules are attracted to the Na⁺ ions, while the partially positive hydrogen atoms in water molecules are attracted to the Cl⁻ ions.

26. How much energy is required to melt 50.0 g of ice at 0 °C? The heat of fusion for water is 6.01 kJ/mol. a. 301 kJ b. 28.9 kJ c. 16.7 kJ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids d. 8.32 kJ e. 156 kJ ANSWER: c 27. Calculate the amount of energy needed to vaporize 125.0 g of water at 100 °C. The heat of vaporization for water is 40.65 kJ/mol. a. 282.0 kJ b. 455.1 kJ c. 169.7 kJ d. 317.2 kJ e. 307.5 kJ ANSWER: a 28. What is the change in energy when 225.0 g of steam at 100 °C condenses to water? The heat of vaporization for water is 40.65 kJ/mol. a. 507.6 kJ b. –258.3 kJ c. 228.7 kJ d. –507.6 kJ e. 258.3 kJ ANSWER: d 29. Calculate the change in energy that occurs when a 25.8 g ice cube at –17.8 °C melts to liquid water at 24.5 °C. The specific heat of ice is 2.02 J/g·°C, the specific heat of liquid water is 4.184 J/g·°C, and the heat of fusion for water is 6.01 kJ/mol. a. 38.2 kJ b. 25.7 kJ c. 12.2 kJ d. 19.9 kJ e. 17.0 kJ ANSWER: c 30. Calculate the change in energy when a 1.25 kg block of solid carbon dioxide (dry ice), at –85.0 °C, sublimes to gaseous carbon dioxide at 25.0 °C. Carbon dioxide sublimes at –78.5 °C. The specific heat capacity for solid carbon dioxide is 0.85 J/g·°C, the specific heat capacity for gaseous carbon dioxide is 0.839 J/g·°C, and the heat Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids of sublimation for carbon dioxide is 25.2 kJ/mol. a. 760 kJ b. 1300 kJ c. 78 kJ d. 830 kJ e. 550 kJ ANSWER: d 31. Which point on a phase diagram represents the conditions under which all three phases (solid, liquid, and gas) coexist in a mixture? a. melting point b. boiling point c. critical point d. triple point e. sublimation point ANSWER: d 32. At which point on a phase diagram does the liquid phase cease to exist, regardless of pressure? a. melting point b. boiling point c. critical point d. triple point e. sublimation point ANSWER: c 33. At which temperature does the vapor pressure of a liquid equal atmospheric pressure? a. melting point b. boiling point c. critical point d. triple point e. sublimation point ANSWER: b 34. On a phase diagram, which phase is typically found at high pressures and low temperatures? a. gas b. liquid c. solid d. plasma e. supercritical fluid Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids ANSWER: c 35. At a given pressure, what is the name of the phase change that involves a solid changing directly into a gas? a. vaporization b. condensation c. fusion d. supercritical e. sublimation ANSWER: e 36. Use the phase diagram to estimate the pressure and temperature for the triple point of ammonia.

a. 1 atm, –78 °C b. 10 atm, 25 °C c. 1 atm, –33 °C d. 0.06 atm, –78 °C e. 15 atm, 30 °C ANSWER: d 37. Use the phase diagram to estimate the normal boiling point of ammonia.

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Chapter 11 Intermolecular Forces, Liquids, and Solids

a. –78 °C b. 25 °C c. –33 °C d. 0 °C e. –50 °C ANSWER: c 38. Under what conditions would a sample of ammonia sublime?

a. At pressures above 0.06 atm and between temperatures of –78 and –33 °C. b. At pressures above 0.06 atm and temperatures above –78 °C. Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids c. At pressures above 10 atm and temperatures above 25 °C. d. At pressures below 0.06 atm and temperatures below –78 °C. e. At pressures below 1 atm and temperatures below –33 °C. ANSWER: d 39. What phase would a sample of ammonia be in at a pressure of 2.5 atm and a temperature of –50 °C?

a. solid b. liquid c. gas d. supercritical fluid e. Ammonia would be at its triple point.

ANSWER: b 40. What phase would a sample of ammonia be in at a pressure of 0.75 atm and a temperature of 20 °C?

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Chapter 11 Intermolecular Forces, Liquids, and Solids

a. solid b. liquid c. gas d. supercritical fluid e. Ammonia would be at its triple point.

ANSWER: c 41. What phase would a sample of ammonia be in at a pressure of 0.06 atm and a temperature of –78 °C?

a. solid b. liquid Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids c. gas d. supercritical fluid e. Ammonia would be at its triple point.

ANSWER: e 42. How much energy is required to completely boil 2.00 L of water from 25.0 °C? The boiling point of water is 100.0 °C. The specific heat capacity for water is 4.184 J/g·°C, the density of water is 1.00 g/mL, and the heat of vaporization for water is 40.65 kJ/mol.

ANSWER: 5,138 kJ 43. Describe the five energy requirements and phase changes required to convert a sample of solid water (ice) below its melting point to gaseous water (steam) at a temperature greater than its boiling point.

ANSWER First, energy is required to heat the ice from its initial temperature to 0 °C, the : melting point of water. Once the ice reaches 0 °C, additional energy is needed to overcome the heat of fusion and convert the solid ice into liquid water at 0 °C. Following this, the liquid water must be heated from 0 °C to 100 °C, which is the boiling point of water. At 100 °C, energy is then required to overcome the heat of vaporization, allowing the liquid water to transition into steam (gaseous phase). Finally, once the water is in the gas phase, further energy is needed to raise the temperature of the steam to a temperature greater than its boiling point.

44. Calculate the change in energy that occurs when an ice cube tray containing 265 g of liquid water at 22.8 °C freezes and ultimately reaches a temperature of –16.5 °C. The specific heat capacity of ice is 2.02 J/g·°C, the specific heat capacity for liquid water is 4.184 J/g·°C, and the heat of fusion is 6.01 kJ/mol.

ANSWER: –122.5 kJ 45. What is a supercritical fluid?

ANSWER A supercritical fluid is a state of matter that occurs at high temperatures and : pressures. In this state, a substance has high kinetic energy like a gas, but the particles are packed close together like a liquid.

46. The vapor pressure of a liquid at 25.0 °C is 100.0 torr, and the heat of vaporization is 40.7 kJ/mol. What is the vapor pressure of the liquid at 50.0 °C? ANSWER: 358 torr 47. The normal boiling point of ethanol is 78.4 °C, and its heat of vaporization is 38.56 kJ/mol. Calculate the vapor pressure of ethanol at 60.0 °C. ANSWER: 367 torr 48. Describe the characteristics of a compound that is considered volatile.

ANSWER: Compounds with a low boiling point and high vapor pressure are considered Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids volatile.

49. Which of the following compounds is likely to have the highest lattice energy? a. lithium nitride b. lithium fluoride c. potassium chloride d. lithium sulfide e. sodium chloride ANSWER: a 50. Which of the following compounds has the highest lattice energy? The melting points are listed with the compounds. a. LiF, 845 °C b. NaCl, 801 °C c. CaO, 2572 °C d. Fe2O3, 1565 °C e. Al2O3, 2072 °C ANSWER: c 51. Which of the following compounds has the lowest lattice energy? The melting points are listed with the compounds. a. LiF, 845 °C b. NaCl, 801 °C c. CaO, 2572 °C d. Fe2O3, 1565 °C e. Al2O3, 2072 °C ANSWER: b 52. Which of the following compounds is likely to have the lowest lattice energy? a. beryllium chloride b. magnesium chloride c. calcium chloride d. strontium chloride e. barium chloride ANSWER: e 53. Which of the following substances is best described as comprising positively and negatively charged ions packed tightly together? a. metals b. molecular compounds Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids c. ionic compounds d. network covalent compounds e. polymers ANSWER: c 54. Which of the following best describes a substance comprised of repeating covalent bonds that extend to two- or three-dimensional lattices? a. metals b. molecular compounds c. ionic compounds d. network covalent compounds e. polymers ANSWER: d 55. Which of the following substances is an example of a naturally occurring polymer? a. sucrose b. sodium chloride c. diamond d. cellulose e. ice ANSWER: d 56. Which of the following substances would best describe a solid at room temperature that is held together by intermolecular forces? a. metals b. molecular compounds c. ionic compounds d. network covalent compounds e. polymers ANSWER: b 57. What is the relationship between lattice energy and melting point for ionic compounds?

ANSWER As the lattice energy increases, the melting point of the ionic compound also : tends to increase.

58. Describe the differences in bonding between molecular compounds and metals.

ANSWER Molecular compounds involve discrete molecules held together by covalent : bonds. Metals share their outer electrons across all atoms in the sample.

59. What is a polymer?

ANSWER: Polymers are molecular compounds containing long chains of repeating units.

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Chapter 11 Intermolecular Forces, Liquids, and Solids 60. How does ionic radii impact the melting point of ionic compounds?

ANSWER Ions with smaller ionic radii lead to stronger electrostatic attractions and higher : melting points, whereas ions with larger ionic radii result in weaker attractions and lower melting points.

61. Which of the following substances is comprised of atoms, molecules, or ions in a well-ordered, repeating structure? a. crystalline solid b. amorphous solid c. geometric solid d. network solid e. polymer solid ANSWER: a 62. Which of the following best describes a solid that contains atoms, molecules, or ions that are NOT in a wellordered, repeating structure? a. crystalline solid b. amorphous solid c. geometric solid d. formless solid e. chaotic solid ANSWER: b 63. What experimental technique allows scientists to determine the structure of crystalline substances? a. mass spectrometry b. magnetic resonance imaging (MRI) c. UV-vis crystallography d. X-ray crystallography e. Nuclear Magnetic Resonance (NMR) spectroscopy ANSWER: d 64. X-rays with a wavelength of 154 pm are used to analyze a crystal. If the first-order diffraction (n = 1) occurs at an angle of 20.0°, what is the distance between the crystal planes? a. 376 pm b. 121 pm c. 175 pm d. 450 pm e. 225 pm ANSWER: e Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids 65. In a diffraction experiment, the distance between crystal planes is 312 pm, and the first-order diffraction (n = 1) occurs at 15.0°. What is the wavelength of the x-rays used? a. 134 pm b. 202 pm c. 161 pm d. 312 pm e. 451 pm ANSWER: c 66. If x-rays of wavelength 71.0 pm produce a second-order diffraction (n = 2) at an angle of 30.0° in a crystal, what is the distance between the planes in the crystal? a. 71.0 pm b. 142 pm c. 284 pm d. 356 pm e. 177 pm ANSWER: b 67. How many atoms occupy a unit cell of a body-centered cubic cell? a. 1 b. 2 c. 4 d. 6 e. 8 ANSWER: b 68. Which close-packed structure positions the third layer of atoms directly above the first layer of atoms? a. cubic-close packed structure b. hexagonal-close packed structure c. centered-close packed structure d. simple-close packed structure e. tetragonal-close packed structure ANSWER: b 69. What is the packing efficiency of a simple cubic cell which contains one atom per unit cell, has an edge length of 2r, and a volume of 8r3? a. 48.6 % b. 52.4 % c. 74.0 % Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids d. 68.0 % e. 64.8 % ANSWER: b 70. What is the density of iron, in g/cm3, when it forms a face-centered cubic structure with an edge length of 365.8 pm? a. 9.12 g/cm3 b. 12.6 g/cm3 c. 8.13 g/cm3 d. 7.58 g/cm3 e. 9.59 g/cm3 ANSWER: d 71. Describe the arrangement of atoms in a unit cell with a simple cubic structure. How many atoms occupy each unit cell?

ANSWER Substances with a simple cubic structure contain an atom at each of the eight : corners. One-eighth of each atom is contained in the unit cell, so the simple cubic structure contains a total of one atom per unit cell.

72. Describe the arrangement of atoms in a unit cell with a face-centered cubic structure. How many atoms occupy each unit cell?

ANSWER Substances with a face-centered cubic structure have an atom at each of the : eight corners and an atom at each of the six faces. One-eighth of each atom occupies each corner in the unit cell, and each face contains half of an atom.

So, the face-centered cubic structure contains a total of four atoms per unit cell.

73. Describe the arrangement of atoms in a unit cell with a body-centered cubic structure. How many atoms are occupied in each unit cell?

ANSWER Substances with a body-centered cubic structure have an atom at each of the : eight corners and contain one atom centered in the unit cell. One-eighth of each atom occupies each corner in the unit cell, and the entire atom in the center is contained inside the unit cell. So, the body-centered cubic structure contains a total of two atoms per unit cell.

74. X-rays with a wavelength of 125.9 pm produce a first-order diffraction (n = 1) at an angle of 12.89° in a crystal. What is the distance between the planes in the crystal?

ANSWER: 282.2 pm 75. What is the edge length (in pm) of tungsten’s body-centered cubic structure? The density of tungsten is 19.3 g/cm3.

ANSWER: 316 pm Copyright Macmillan Learning. Powered by Cognero.

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Chapter 11 Intermolecular Forces, Liquids, and Solids

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Chapter 10 Gases 1. A car's tire has a maximum pressure rating of 35 psi. What is the value of the pressure rating in torr? a. 1800 torr b. 0.68 torr c. 3500 torr d. 0.34 torr e. 770 torr ANSWER: a 2. The planet Venus has an atmospheric pressure of 1350 psi. What is the atmosphere pressure rating in bar? a. 1.78 bar b. 22.4 bar c. 14.7 par d. 93.0 bar e. 106 bar ANSWER: d 3. What is the pressure of a gas connected to a mercury manometer if the atmospheric pressure is 758 torr and the mercury level open to the gas is 21 mm higher than the mercury open to the atmosphere? a. 758 torr b. 779 torr c. 800 torr d. 737 torr e. 716 torr ANSWER: d 4. Which of the following is not a property of an ideal gas? a. The volume of the particles is less than the total volume the gas occupies. b. The particles have no attraction to each other. c. The collisions between particles are elastic. d. The velocity of the particles does not change as they pass near each other. e. The total momentum of the gas particles decreases when the particles collide.

ANSWER: e 5. Which of the following is the highest pressure? a. 865 torr b. 15.7 psi c. 1.15 atm d. 775 mm Hg e. 125 kPa ANSWER: e Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases 6. Calculate the pressure in kilopascals exerted by a column of mercury with a height of 1.25 m. The density of mercury is 1.36×104 kg/m3. a. 167 kPa b. 107 kPa c. 125 kPa d. 89.9 kPa e. 91.2 kPa ANSWER: a 7. A mercury manometer is open to the atmosphere on one side, and connected to a gas on the other. Based on the figure, if the atmospheric pressure is 765 torr, what is the pressure of the gas?

a. 553 torr b. 977 torr c. 659 torr d. 871 torr e. 825 torr ANSWER: b 8. What is the range of pressure that water exerts on a deep-sea fish that inhabits a depth of 1,000–4,000 m? The density of water is 1.00×103 kg/m3. Report your answers in atmospheres and with three significant figures.

ANSWER: The range of pressure that water exerts is approximately 96.8 atm to 387 atm.

9. A mercury manometer is open to the atmosphere on one side, and connected to a gas on the other. Based on the figure, if the atmospheric pressure is 772 torr, what is the pressure of the gas?

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Chapter 10 Gases

ANSWER: 465 torr 10. The atmosphere of Mars is much thinner and colder than Earth's, having a pressure of 0.088 psi. How many times greater is Earth’s atmospheric pressure of 760 torr than Mars’ atmospheric pressure?

ANSWER Earth's atmospheric pressure is approximately 170 times greater than Mars' : atmospheric pressure.

11. Which gas law describes the relationship between a gas's pressure, volume, and temperature, assuming the number of moles remains constant? a. Boyle's law b. Charles's law c. Avogadro's law d. Gay-Lussac's law e. combined gas law ANSWER: e 12. If the volume of a gas is halved while keeping the temperature and number of moles constant, what happens to its pressure? a. it doubles b. it halves c. it quadruples d. it remains the same e. it decreases by a factor of four ANSWER: a 13. A gas sample occupies a volume of 3.0 L at 25 °C and 1.0 atm. If the temperature is increased to 50 °C while keeping the pressure constant, what is the new volume of the gas? a. 1.5 L Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases b. 2.8 L c. 3.3 L d. 6.0 L e. 9.0 L ANSWER: c 14. A gas occupies a volume of 2.0 L at a pressure of 1.0 atm and a temperature of 273 K. If the volume is increased to 4.0 L and the temperature is increased to 546 K, what is the new pressure of the gas, assuming the number of moles remains constant? a. 0.50 atm b. 1.0 atm c. 2.0 atm d. 4.0 atm e. 8.0 atm ANSWER: b 15. A 5.0 L container holds a gas at a pressure of 2.0 atm and a temperature of 27 °C. How many moles of gas are in the container? a. 4.5 mol b. 1.7 mol c. 0.23 mol d. 0.41 mol e. 2.7 mol ANSWER: d 16. A 10.0 L gas cylinder holds 2.75 moles of helium gas at a temperature of 25.0 °C. What is the pressure of the gas in the container? a. 4.51 atm b. 6.73 atm c. 0.628 atm d. 2.39 atm e. 1.75 atm ANSWER: b 17. According to Avogadro's law, at constant temperature and pressure, the volume of a gas is directly proportional to its a. pressure. b. temperature. c. mass.

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Chapter 10 Gases d. number of moles. e. density.

ANSWER: d 18. A sample of nitrogen occupies a volume of 3.0 L at a temperature of 20 °C. If the temperature is increased to 60 °C while keeping the pressure constant, what will be the new volume of the gas? a. 2.1 L b. 3.4 L c. 5.6 L d. 6.1 L e. 9.0 L ANSWER: b 19. A helium balloon at sea level, with an atmospheric pressure of 14.7 psi, has a volume of 2.45 L. What would the volume of the balloon be if the balloon reached an altitude of 10,000 ft with an atmospheric pressure of 10.7 psi? a. 2.01 L b. 2.67 L c. 3.37 L d. 0.298 L e. 64.2 L ANSWER: c 20. A 5.0 L container holds a gas at a pressure of 2.0 atm and a temperature of 27 °C. How many moles of gas are in the container? a. 4.5 mol b. 1.7 mol c. 0.23 mol d. 0.41 mol e. 2.7 mol ANSWER: d 21. How many grams of nitrogen are produced in a reaction if as a result of its production a 750.0 mL reaction vessel has a pressure of 17.3 psi at 26.2 °C? ANSWER: 1.01 g 22. A 3.00 L balloon is filled with helium at a pressure of 1.00 atm and a temperature of 273.0 K. If the balloon is taken to a place where the pressure is 0.80 atm, and the temperature is 300.0 K, what will be the new volume of the balloon?

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Chapter 10 Gases ANSWER: 4.11 L 23. Aerosol spray cans have a warning label that states: “Never store aerosol cans at temperatures above 120 °F”. Why do the manufacturers place this warning label on the cans?

ANSWER Manufacturers place this warning label on aerosol spray cans because, : according to Charles's law, as the temperature increases, the volume (or pressure if the volume is constant) of the gas inside the can increases, which can lead to the can exploding at high temperatures.

24. What is the relationship between the volume and pressure of a gas at constant temperature?

ANSWER The relationship between the volume and pressure of a gas at constant : temperature is described by Boyle's law, which states that the volume of a gas is inversely proportional to its pressure.

25. What is the relationship between the volume and temperature of a gas at constant pressure?

ANSWER The relationship between the volume and temperature of a gas at constant : pressure is described by Charles's law, which states that the volume of a gas is directly proportional to its absolute temperature (measured in Kelvin).

26. What is the relationship between the volume and number of moles of a gas at constant temperature and pressure?

ANSWER The relationship between the volume and the number of moles of a gas at : constant temperature and pressure is described by Avogadro's law, which states that the volume of a gas is directly proportional to the number of moles of gas present.

27. Which of the following gases would have the highest density at atmospheric pressure and room temperature? a. Ar b. H2 c. CO2 d. N2 e. O2 ANSWER: c 28. A 12.00 g sample of an unknown gas occupies a volume of 2.240 L at 25.0 °C and 14.89 psi. What is the molar mass of the gas? a. 157.4 g/mol b. 79.0 g/mol c. 129.3 g/mol d. 89.9 g/mol e. 64.65 g/mol Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases ANSWER: c 29. Which of the following gases would have the lowest density at atmospheric pressure and room temperature? a. Ne b. He c. CH4 d. NO2 e. SO2 ANSWER: b 30. A 2.50 g sample of an unknown gas occupies a volume of 2.20 L at 365 K and 1.00 atm. What is the molar mass of the gas? a. 34.0 g/mol b. 68.1 g/mol c. 103 g/mol d. 75.2 g/mol e. 19.5 g/mol ANSWER: a 31. What is the density of oxygen at 1.00 atm and 25 °C? a. 1.31 g/L b. 1.43 g/L c. 1.14 g/L d. 1.00 g/L e. 1.25 g/L ANSWER: a 32. A 1.244 kg sample of a gas in a 24.0 L container has a pressure of 291.42 psi at 25 °C. What is the molar mass of the gas? a. 105 g/mol b. 36.5 g/mol c. 47.9 g/mol d. 64.0 g/mol e. 58.2 g/mol ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases 33. What is the density of carbon dioxide at 772 torr and 12 °C? a. 1.27 g/L b. 1.91 g/L c. 0.987 g/L d. 2.14 g/L e. 1.64 g/L ANSWER: b 34. What causes paper lanterns with lit candles to ascend into the air?

ANSWER The heat generated by the lit candle warms the air inside the lantern, causing it : to expand and become less dense than the cooler air outside the lantern.

35. Rank the following gases in order of increasing density at room temperature and atmospheric pressure: Xe, CO2, Cl2, F2, N2.

ANSWER The gases ranked in order of increasing density at room temperature and : atmospheric pressure are: N2 < F2 < CO2 < Cl2 < Xe 36. A 15.00 g sample of an unknown gas is placed in a 2.00 L container at a temperature of 298 K and a pressure of 44.1 psi. Calculate the molar mass of the gas.

ANSWER: The molar mass of the gas is 61.2 g/mol.

37. A gas cylinder contains 60.0 g of a certain gas at a pressure of 4.05 bar and a temperature of 273 K in a volume of 5.00 L. Determine the molar mass of the gas.

ANSWER: The molar mass of the gas is 67.2 g/mol.

38. A gas mixture contains 1.25 moles of oxygen and 3.75 moles of nitrogen in a 10.0 L container at a certain temperature. If the total pressure of the gas mixture is 4.00 atm, what is the partial pressure of nitrogen? a. 1.25 atm b. 3.00 atm c. 2.50 atm d. 1.00 atm e. 2.75 atm ANSWER: b 39. A 4.00 L container holds a mixture of He and Ar gases at 400 K. The mixture contains 1.50 moles of He and 2.50 moles of Ar. If the total pressure in the container is 6.00 atm, what is the mole fraction of helium? a. 0.600 b. 0.625 c. 0.500 d. 0.425 e. 0.375 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases ANSWER: e 40. A mixture of gases contained in a 10.0-L gas cylinder has a pressure of 205 psi at 295 K. If the cylinder contains 15.5 g of nitrogen (N2), what is the partial pressure of nitrogen? a. 14.0 atm b. 0.996 atm c. 1.34 atm d. 5.62 atm e. 1.75 atm ANSWER: c 41. A 15.0-L gas cylinder holds a mixture of gases. The pressure inside the cylinder is 10.2 atm at 298 K. What is the mole fraction of argon in the mixture if there is 15.5 g of argon inside the cylinder? a. 0.0621 b. 0.120 c. 0.362 d. 0.0482 e. 0.758 ANSWER: a 42. A container with a total pressure of 150 psi contains one mole of each of the following gases: Xe, Ne, He, and Ar. Which gas would have the highest partial pressure?

ANSWER Since the mole fraction is the same, all gases (Xe, Ne, He, and Ar) would have : the same partial pressure.

43. What is the mole fraction of hydrogen and oxygen in a sample containing 12.5 g of H2 and 17.5 g of O2 at a pressure of 25.0 psi?

ANSWER Therefore, the mole fraction of hydrogen is 0.919, and the mole fraction of : oxygen is 0.0810.

44. What is the partial pressure of hydrogen and nitrogen in a reaction vessel containing 35.1 g of H2 and 56.9 g of N2 at a pressure of 155 psi?

ANSWER The partial pressure of hydrogen is 139 psi, and the partial pressure of nitrogen : is 16.0 psi.

45. Which of the following gases would have the highest velocity at 298 K? a. He b. Ne c. CO2 d. NH3 e. Xe Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases ANSWER: a 46. Which of the following gases would have the lowest velocity at 298 K? a. H2 b. NO2 c. CH4 d. CO e. CH3CH3 ANSWER: b 47. What is the ump for oxygen at 375 K? a. 496 m/s b. 441 m/s c. 140 m/s d. 129 m/s e. 292 m/s ANSWER: b 48. Which of the following gases has the lowest diffusion rate? a. He b. CO2 c. NH3 d. Xe e. CH4 ANSWER: d 49. What is the term that describes the rate at which a gas escapes from a container? a. most probable speed b. diffusion c. effusion d. compressibility factor e. dispersion ANSWER: c 50. Which of the following gases has the highest effusion rate? a. H2 b. CO2 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases c. Ar d. CO e. He ANSWER: a 51. How does the kinetic theory of gases explain the pressure exerted by a gas?

ANSWER The pressure exerted by a gas is caused by the collisions of gas molecules with : the walls of the container. Each collision exerts a force on the wall, and the collective impact of these collisions per unit area results in the pressure of the gas.

52. What is the relationship between the average speed of gas particles and molar mass, according to the kinetic theory of gases?

ANSWER According to the kinetic theory of gases, the average speed of gas particles is : inversely related to the square root of their molar mass. Thus, lighter gas particles move faster on average than heavier gas particles at the same temperature.

53. According to the kinetic theory of gases, why does pressure increase with temperature?

ANSWER Pressure increases with temperature because the average kinetic energy of the : gas molecules increases as the temperature rises. When the temperature of a gas increases, the molecules move faster. This increased molecular speed results in more frequent and more forceful collisions with the walls of the container, causing an increase in pressure.

54. Glove boxes are sealed chambers used for air-sensitive chemical reactions. Which gas, nitrogen or argon, would be a better choice to fill the chamber to minimize leakage? Considering Graham's law, calculate the relative effusion rate of the chosen gas compared to the other gas.

ANSWER Argon would be a better choice to fill the glove box to minimize leakage : because it effuses more slowly than nitrogen. The relative effusion rate of argon compared to nitrogen is 0.8375.

55. Hydrogen can be produced from the electrolysis of water according to the equation below. What would the partial pressure of hydrogen be in a 1.25 L container at 26.5 °C if 25.2 g of water reacted?

2H2O(l) → 2H2(g) + O2(g) a. 7.85 atm b. 18.2 atm c. 27.5 atm d. 12.9 atm e. 13.0 atm ANSWER: c 56. Ammonia is produced from hydrogen and nitrogen according to the following equation. A 25.0-L reaction Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases vessel containing only ammonia has a pressure of 1.65 atm at 27.2 °C. How many grams of nitrogen reacted to form the amount of ammonia in the reaction vessel?

3H2(g) + N2(g) → 2NH3(g) a. 17.8 g b. 23.5 g c. 27.5 g d. 106 g e. 62.9 g ANSWER: b 57. How many grams of oxygen are required to react completely with 2.5 L of hydrogen at 25 °C and 1.05 atm to produce water?

2H2(g) + O2(g) → 2H2O(g) a. 0.0536 g b. 3.44 g c. 1.72 g d. 0.860 g e. 1.13 g ANSWER: c 58. The refrigerant, CCl2F2, can be produced according to the following equation. What pressure of CCl2F2 is produced if 9.08 g of HF reacts with excess CCl4 in a 2.0-L container at 40.5 °C? CCl4(g) + 2HF(g) → CCl2F2(g) + 2HCl(g) a. 1.47 atm b. 3.07 atm c. 1.92 atm d. 5.86 atm e. 2.93 atm ANSWER: e 59. Sodium reacts with water to produce hydrogen gas according to the following reaction. If 5.68 g of sodium is added to 100.0 mL of water, what volume of hydrogen gas is produced at STP? 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) a. 1.25 L b. 2.35 L c. 0.857 L d. 2.77 L e. 1.69 L ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases 60. The refrigerant, CCl2F2, can be produced according to the following equation. What pressure of HF is required to produce 1.25 mol of CCl2F2 in a 5.0-L container at 75.0 °C? Assume there is excess CCl4 present.

CCl4(g) + 2HF(g) → CCl2F2(g) + 2HCl(g) a. 14.3 atm b. 7.16 atm c. 28.6 atm d. 12.5 atm e. 2.08 atm ANSWER: a 61. What volume of carbon dioxide is produced after 12.5 g of CaO is produced from heating CaCO3 at 775 K and 762 torr?

CaCO3(s) → CaO(s) + CO2(g) a. 7.11 L b. 14.2 L c. 12.5 L d. 6.98 L e. 28.4 L ANSWER: b 62. Methane is produced from hydrogen and carbon monoxide according to the following equation. A 20.0-L reaction vessel contains hydrogen and carbon monoxide at 3.50 atm at 85.0 °C. The partial pressure of hydrogen was measured to be 2.10 atm. What is the limiting reactant and theoretical yield of methane for this reaction?

CO(g) + 3H2(g) → CH4(g) + H2O(g) ANSWER: The limiting reactant is hydrogen, and the theoretical yield of methane is 7.65 grams.

63. Ammonia is produced from hydrogen and nitrogen according to the following equation. A 15.0-L reaction vessel contains hydrogen and nitrogen at 2.75 atm at 57.8 °C. The partial pressure of hydrogen was measured to be 1.87 atm. What is the limiting reactant and theoretical yield of ammonia for this reaction?

3H2(g) + N2(g) → 2NH3(g) ANSWER: The limiting reactant is hydrogen, and the theoretical yield of ammonia is 11.7 grams.

64. How many grams of FeS are required to produce hydrogen sulfide with a pressure of 35.6 psi in a 15.0-L container at 55.5 °C? Assume excess hydrochloric acid is present.

FeS(s) + 2 HCl(aq) → FeCl2(aq) + H2S(g) ANSWER: The mass of FeS required to produce that amount of hydrogen sulfide is 118.46 grams.

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Chapter 10 Gases 65. Which of the following is a primary reason that real gases deviate from ideal gas behavior at high pressures? a. The temperature of the gas decreases. b. The volume of the gas particles becomes significant compared to the container volume. c. The gas particles move faster. d. The intermolecular forces become negligible. e. The gas particles undergo elastic collisions.

ANSWER: b 66. Why do real gases deviate from ideal gas behavior at low temperatures? a. Gas particles are in constant, random motion. b. Gas particles have different volumes. c. Intermolecular forces become significant. d. Gas particles undergo perfectly elastic collisions. e. The pressure of the gas increases.

ANSWER: c 67. Under which conditions are deviations from the ideal gas law most likely to be observed? a. low pressure and high temperature b. high pressure and high temperature c. low pressure and low temperature d. high pressure and low temperature e. standard temperature and pressure ANSWER: d 68. Which statement best explains why real gases do not always behave ideally? a. Real gases are composed of particles that have a mass. b. Real gases have negligible intermolecular attractions. c. Real gases have elastic collisions with no energy loss. d. Real gases have particles with significant volume and intermolecular forces. e. Real gases follow the Ideal Gas Law at all temperatures and pressures.

ANSWER: e 69. What happens to real gases as they are compressed to high pressures? a. The kinetic energy of the gas particles becomes negligible. b. The volume occupied by the gas particles becomes negligible c. The collisions between gas particles are perfectly elastic. d. The volume occupied by the gas particles becomes significant. e. The gas particles move independently without any interaction.

ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 10 Gases 70. Which of the following gases is most likely to have a compressibility factor of less than one? a. N2 b. Ne c. NH3 d. CH4 e. C2H6 ANSWER: c 71. Which of the following gases has the largest a parameter in the van der Waals equation? a. CH4 b. CO2 c. HCl d. H2O e. F2 ANSWER: d 72. Which of the following gases has the largest b parameter in the van der Waals equation? a. He b. Kr c. Ne d. Ar e. Cl2 ANSWER: e 73. A 15.0-L cylinder contains 4.50 mols of HF at 27.8 °C. What is the pressure of the cylinder using the van der Waals equation? For HF, a = 9.431 atm·L2/mol2 and b = 0.0739 L/mol. a. 6.73 atm b. 11.0 atm c. 3.57 atm d. 9.32 atm e. 4.25 atm ANSWER: a 74. Use the van der Waals equation to determine the volume of a 5.65 mol sample of ammonia gas at 1250 torr and 375 K. For ammonia, a = 4.17 atm·L2/mol2 and b = 0.0371 L/mol.

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Chapter 10 Gases a. 78.2 L b. 110 L c. 57.3 L d. 98.2 L e. 105 L ANSWER: a 75. At 375 K and a pressure of 20.0 atm, a gas has a compressibility factor of 1.3. Calculate the volume occupied by one mole of this gas. Determine whether the observed volume is higher or lower than that predicted by the ideal gas law.

ANSWER The observed volume is 2.00 L, and is higher than the volume predicted by the : Ideal Gas Law which is 1.54 L. This is because the compressibility factor is greater than 1.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models 1. How many electron sets are around the carbon atom in hydrogen cyanide, HCN? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: b 2. How many electron sets are around the N atom in the nitrite ion?

a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: c 3. How many electron sets are around the C atom in the carbonate ion, CO32– ? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: c 4. How many electron sets are around a central atom with a trigonal pyramidal molecular geometry? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: d 5. What is the electron geometry and the molecular geometry for H2S? a. electron geometry = trigonal planar; molecular geometry = trigonal planar b. electron geometry = linear; molecular geometry = linear c. electron geometry = tetrahedral; molecular geometry = trigonal pyramidal Copyright Macmillan Learning. Powered by Cognero.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models d. electron geometry = tetrahedral; molecular geometry = tetrahedral e. electron geometry = tetrahedral; molecular geometry =bent ANSWER: e 6. What is the molecular geometry associated with the C atom in methanol, CH3OH? a. trigonal planar b. linear c. trigonal pyramidal d. tetrahedral e. bent ANSWER: d 7. What is the molecular geometry of the carbonate ion, CO32– ? a. linear b. trigonal planar c. tetrahedral d. trigonal pyramidal e. bent ANSWER: b 8. What is the molecular geometry for CH2Cl2? a. trigonal planar b. linear c. trigonal pyramidal d. tetrahedral e. bent ANSWER: d 9. What is the electron geometry and the molecular geometry for PH3? a. electron geometry = trigonal planar; molecular geometry = trigonal planar b. electron geometry = linear; molecular geometry = linear c. electron geometry = tetrahedral; molecular geometry = trigonal pyramidal d. electron geometry = tetrahedral; molecular geometry = tetrahedral e. electron geometry = tetrahedral; molecular geometry =bent ANSWER: c 10. Which molecular shape corresponds to a molecule with a central atom with two lone pairs of electrons and Copyright Macmillan Learning. Powered by Cognero.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models an electronic geometry of trigonal bipyramidal? a. trigonal planar b. linear c. trigonal pyramidal d. t-shape e. bent ANSWER: d 11. What are the molecular geometries for the three atoms marked as a, b, and c in the following molecule?

a. a = tetrahedral, b = trigonal pyramidal , and c = trigonal pyramidal b. a = trigonal pyramidal , b = trigonal planar, and c = tetrahedral c. a = tetrahedral, b = trigonal planar, and c = trigonal pyramidal d. a = tetrahedral, b = bent, and c = trigonal pyramidal e. a = tetrahedral, b = trigonal planar, and c = trigonal planar ANSWER: c 12. What is the bond angle for a molecule with a trigonal planar molecular geometry? a. 180° b. 145° c. 120° d. 109.5° e. 90° ANSWER: c 13. Which model predicts the arrangement of electron sets around a central atom by maximizing the distance between them? a. Lewis structure b. molecular orbital theory c. valence shell electron pair repulsion d. hybridization e. atomic theory ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models 14. Which term describes the spatial arrangement of atoms within a molecule? a. electronic geometry b. molecular geometry c. valence bond theory d. hybridization e. molecular symmetry ANSWER: b 15. Describe the bonding and electron set arrangement that results in a molecule with a T-shape molecular geometry.

ANSWER A T-shaped molecular geometry is observed in molecules with three bonded : atoms and two lone pairs of electrons on the central atom.

16. How does the concept of molecular geometry differ from electronic geometry?

ANSWER Molecular geometry refers to the arrangement of atoms within a molecule, : while electronic geometry describes the arrangement of electron sets around a central atom.

17. How does valence shell electron pair repulsion model (VSEPR) predict the electron geometry of molecules?

ANSWER VSEPR theory predicts the electronic geometry of a molecule by assuming that : electron sets adopt arrangements that minimize the repulsion between them, thereby maximizing the distance between electron sets.

18. Draw the Lewis structure and determine the molecular geometry of HCN. ANSWER: The molecular geometry for HCN is linear.

19. Describe the molecular geometry of a molecule with the formula CCl4 and explain how it is determined using the VSEPR theory.

ANSWER The molecular geometry of CCl4 is tetrahedral, determined by the arrangement : of four bonding electron pairs around the central carbon atom, minimizing repulsions to achieve the most stable geometry.

20. What are the axial and equatorial bond angles for a molecule with a trigonal bipyramidal molecular geometry?

ANSWER For a molecule with a trigonal bipyramidal molecular geometry, the atoms in : axial positions are 90° from the equatorial atoms and 180° from the other axial atom. The atoms in equatorial positions are 90° from the axial atoms and 120∘ from the other equatorial atoms.

21. A nonpolar covalent bond is characterized by Copyright Macmillan Learning. Powered by Cognero.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models a. equal sharing of electrons between atoms. b. unequal sharing of electrons between atoms. c. complete transfer of electrons from one atom to another. d. no sharing of electrons between atoms. e. random sharing of electrons between atoms.

ANSWER: a 22. Which of the following molecules is polar? a. CH4 b. CCl4 c. CO2 d. CH3OH e. BF3 ANSWER: d 23. What is the range of electronegativity difference between atoms for a bond to be classified as nonpolar covalent? a. less than 0.5 b. 0.5 to 2.0 c. greater than 2.0 d. equal to 2.0 e. equal to 0.5 ANSWER: a 24. A polar covalent bond is characterized by a. equal sharing of electrons between atoms. b. unequal sharing of electrons between atoms. c. complete transfer of electrons from one atom to another. d. no sharing of electrons between atoms. e. random sharing of electrons between atoms.

ANSWER: b 25. Which of the following molecules is nonpolar? a. HF b. CF4 c. CH2O d. CH3CH2OH e. NH2CH3 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models ANSWER: b 26. Which of the following bonds is ionic? a. H–F b. O–H c. C–H d. Na–F e. N–H ANSWER: d 27. Which of the following bonds is nonpolar? a. H–F b. O–H c. C–H d. Na–F e. N–H ANSWER: c 28. Which of the following molecules does not have a net dipole moment? a. CH2Cl2 b. CCl4 c. NH3 d. CO e. CH3F ANSWER: b 29. Which of the following molecules has a net dipole moment? a. CH4 b. CCl4 c. NH3 d. CO2 e. BF3 ANSWER: c 30. Rank the following elements in order of increasing electronegativity: F, C, H, Ca. a. Ca < H < C < F b. H < Ca < C < F c. C < H < Ca < F Copyright Macmillan Learning. Powered by Cognero.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models d. F < C < H < Ca e. F < H < C < Ca ANSWER: a 31. Which of the following elements has the highest electronegativity? a. Mg b. Sn c. P d. N e. S ANSWER: d 32. Which of the following elements has the lowest electronegativity? a. C b. Cl c. B d. Al e. Ca ANSWER: e 33. Which of the following bonds is polar covalent? a. Na–O b. C–H c. C–C d. O–H e. F–F ANSWER: d 34. How many of the following bonds would result in a partial positive charge on the C atom: C–H, C–O, C–S, C–Cl, C–N? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: c 35. Which of the following statements correctly describes the polarity of the following molecule?

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models

a. The molecule is nonpolar. b.The molecule has a polar bond between the C and Cl atoms. Carbon has a partial positive charge and chlorine has a partial negative charge. c. The molecule has a polar bond between the C and Cl atoms. Carbon has a partial negative charge and chlorine has a partial positive charge. d.The molecule has polar bonds between the C and the H atoms. Carbon has a partial positive charge, while the H atoms contain partial negative charges. e. The molecule has polar bonds between the C and the H atoms. Carbon has a partial negative charge, while the H atoms contain partial positive charges.

ANSWER: b 36. What is the electron geometry and molecular geometry of XeF4? Is XeF4 polar or nonpolar?

ANSWER Xenon tetrafluoride has an octahedral electron geometry and square planar : molecular geometry. The square planar geometry of the xenon-fluorine bonds causes the dipole moments to cancel, resulting in a nonpolar molecule.

37. How are bonds categorized into nonpolar covalent, polar covalent, or ionic based on electronegativity differences between atoms?

ANSWER The electronegativity difference between atoms determines the type of bond : formed. Bonds are classified into three categories based on this difference: nonpolar covalent (less than 0.5), polar covalent (0.5 to 2.0), and ionic (greater than 2.0).

38. Is CH3Cl a polar or nonpolar molecule? If polar, describe the arrangement of partial charges and dipole moment.

ANSWER CH3Cl is a polar molecule. There is a partial negative charge on the Cl atom, : and a partial positive charge on the central C atom. There is a net dipole moment from the central C atom to the chlorine atom.

39. Is ammonia, NH3, a polar or nonpolar molecule? If polar, describe the arrangement of partial charges and dipole moment.

ANSWER NH3 is a polar molecule. There is a partial negative charge on the N atom, and : a partial positive charge on each of the H atoms. There is a net dipole moment toward the central N atom.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models 40. Why is carbon tetrachloride a nonpolar molecule despite the carbon-chlorine bond being polar?

ANSWER Carbon tetrachloride has a tetrahedral molecular geometry. The dipole : moments associated with the carbon-chlorine bond cancel out because they pull in opposite directions, resulting in no net dipole moment.

41. How many sigma bonds and pi bonds are found in a carbon dioxide, CO2, molecule? a. two sigma bonds and zero pi bonds b. zero sigma bonds and two pi bonds c. one sigma bond and one pi bond d. two sigma bonds and one pi bond e. two sigma bonds and two pi bonds ANSWER: e 42. Which type of bond results from head-on overlap of atomic orbitals? a. sigma bond b. pi bond c. delta bond d. phi bond e. gamma bond ANSWER: a 43. What is the hybridization for the three atoms marked as a, b, and c in the following molecule?

a. a = sp3, b = sp2, and c = sp3 b. a = sp3, b = sp2, and c = sp2 c. a = sp2, b = sp2, and c = sp3 d. a = sp2, b = sp3, and c = sp2 e. a = sp3, b = sp3, and c = sp2 ANSWER: a 44. What is the hybridization of nitrogen in ammonia, NH3? a. sp b. sp2 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models c. sp3 d. sp3d e. sp3d2 ANSWER: c 45. What is the hybridization of carbon in HCN? a. sp b. sp2 c. sp3 d. sp3d e. sp3d2 ANSWER: a 46. What is the hybridization of an atom with a tetrahedral electron geometry? a. sp b. sp2 c. sp3 d. sp3d e. sp3d2 ANSWER: c 47. What orbitals in the sulfur atom hybridize to form the H2S molecule? a. 2s and 2p b. 3s and 3p c. 3p and 3d d. 3s, 3p, and 3d e. 2p and 3p ANSWER: b 48. What orbitals in the sulfur atom hybridize to form the SF6 molecule? a. 2s and 2p b. 3s and 3p c. 3p and 3d d. 3s, 3p, and 3d e. 2p and 3p Copyright Macmillan Learning. Powered by Cognero.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models ANSWER: d 49. What electron geometry would result from a molecule with an sp2 hybridization? a. linear b. trigonal planar c. tetrahedral d. trigonal bipyramidal e. octahedral ANSWER: b 50. What is the bond angle for a molecule that has a central atom with an sp hybridization? a. 180° b. 145° c. 120° d. 109.5° e. 90° ANSWER: a 51. What electron geometry would result from a molecule with an sp3d2 hybridization? a. linear b. trigonal planar c. tetrahedral d. trigonal bipyramidal e. octahedral ANSWER: e 52. What is the valence overlap model, and how does it explain the formation of covalent bonds?

ANSWER The valence overlap model postulates that covalent bonds form when the : valence orbitals of two atoms overlap with each other, allowing for the pairing of valence electrons with opposite spins from each atom. These overlapping half-filled orbitals lead to a shared electron density between the atoms, constituting a covalent bond. The greater the overlap, the stronger the bond formed.

53. Describe the concept of hybridization in the context of chemical bonding.

ANSWER Hybridization is a theoretical model that explains the bonding in molecules by : mixing different types of atomic orbitals (s, p, d, etc.) on the same atom to produce new orbitals of equal energy called hybrid orbitals. These hybrid orbitals have shapes and orientations that are more suitable for the formation of chemical bonds.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models 54. Describe the electronic geometry, molecular geometry, and hybridization around the central atom in PH3.

ANSWER The phosphorus atom in PH3 has four electron groups around it: three bonding : pairs and one lone pair. This arrangement results in a tetrahedral electronic geometry and a trigonal pyramidal molecular geometry. To accommodate the four electron groups, phosphorus hybridizes one s orbital and three p orbitals to form four equivalent sp3 hybrid orbitals used for bonding with the hydrogen atoms.

55. What are the differences between sigma (σ) and pi (π) bonds in terms of their formation and properties?

ANSWER The initial bond formed between atoms is a head-on overlap of atomic orbitals, : resulting in a sigma (σ) bond. The subsequent bonds formed between atoms, however, arise from the side-by-side interaction of unhybridized p orbitals.

These bonds are known as pi (π) bonds.

56. Why can sulfur form six bonds with a sp3d2 hybridization, but oxygen, which is in the same group, cannot?

ANSWER Sulfur's electron configuration includes empty 3d orbitals and 3s and 3p : orbitals. When sulfur forms compounds, these empty d orbitals can participate in hybridization, leading to the sp3d2 hybridization scheme. This hybridization allows sulfur to form six bonds by combining one s orbital, three p orbitals, and two d orbitals.

Oxygen has no 2d orbitals; therefore, oxygen cannot utilize d orbitals in hybridization. As a result, oxygen typically forms bonds using only its 2s and 2p orbitals.

57. What is sp3 hybridization, and which molecular geometry does it typically correspond to?

ANSWER sp3 hybridization occurs when one s orbital and three p orbitals on the same : atom mix to form four equivalent sp3 hybrid orbitals. This hybridization is typically seen in atoms with tetrahedral geometry, where the bond angles are approximately 109.5º. An example of a molecule with sp3 hybridization is methane (CH4).

58. What type of hybridization is expected for the central atom of a molecule with a trigonal planar molecular geometry? Briefly explain the reasoning for the answer.

ANSWER A molecule with a trigonal planar geometry typically has sp2 hybridization in : the central atom. This hybridization involves combing one s orbital with two p orbitals to form three sp2 hybrid orbitals arranged in a plane 120º apart. This arrangement allows for adequate overlap and bonding with three other atoms or groups in a planar configuration.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models 59. Describe the hybridization for each carbon atom with respect to its effect on the types of bonds formed and the geometry of ethene, C2H4.

ANSWER In ethene (C2H4), each carbon atom undergoes sp2 hybridization. This means : that one s orbital and two p orbitals hybridize to form three sp2 hybrid orbitals, which lie in a plane with 120º angles between them. The remaining p orbital, which is unhybridized, lies perpendicular to this plane. Each carbon uses its three sp2 orbitals to form sigma bonds: two with hydrogen atoms and one with the other carbon atom. The unhybridized p orbitals from each carbon overlap side-to-side to form a pi bond, which, combined with the sigma bond between the carbon atoms, results in a double bond.

60. What is sp2 hybridization, and which molecular geometry does it typically correspond to?

ANSWER sp2 hybridization occurs when one s orbital and two p orbitals on the same : atom hybridize to form three equivalent sp2 hybrid orbitals. This hybridization is typically seen in atoms with trigonal planar geometry, where the bond angles are approximately 120º.

61. What is the significance of the bond order in molecular orbital theory? a. It indicates the number of atoms in a molecule. b. It represents the strength of the bonds in a molecule. c. It determines the molecular geometry of a molecule. d. It reflects the electron configuration of a molecule. e. It predicts the hybridization of a molecule.

ANSWER: b 62. In molecular orbital theory, the higher energy orbitals of bonding atoms are called the a. bonding orbitals. b. antibonding orbitals. c. nodes. d. bond orders. e. atomic orbitals.

ANSWER: b 63. What is the bond order for hydrogen, H2?

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models

a. 0 b. 0.5 c. 1 d. 1.5 e. 2 ANSWER: c 64. In molecular orbital theory, the lower energy orbitals of bonding atoms are called the a. bonding orbitals. b. antibonding orbitals. c. nodes. d. bond orders. e. atomic orbitals.

ANSWER: a 65. Which of the following bond order values would predict the strongest bond? a. 0 b. 0.5 c. 1 d. 1.5 e. 2 ANSWER: e 66. Which of the following best describes a bonding molecular orbital? a. A molecular orbital formed by the constructive interference of atomic orbitals. b. A molecular orbital formed by the destructive interference of atomic orbitals. c. A molecular orbital with a higher energy level than the atomic orbitals. d. A molecular orbital that contains a node. e. A molecular orbital with the same energy level as the atomic orbitals.

ANSWER: a 67. In molecular orbital theory, anitbonding molecular orbitals are formed by Copyright Macmillan Learning. Powered by Cognero.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models a. hybridization of s and p orbitals. b. hybridization of s and d orbitals. c. nonoverlapping of atomic orbitals. d. constructive interference of atomic orbitals. e. destructive interference of atomic orbitals.

ANSWER: e 68. Which of the following best describes an antibonding molecular orbital? a. A molecular orbital formed by the constructive interference of atomic orbitals. b. A molecular orbital formed by the hybridization of atomic orbitals. c. A molecular orbital with a higher energy level than the atomic orbitals. d. A molecular orbital with a lower energy level than the atomic orbitals. e. A molecular orbital with the same energy level as the atomic orbitals.

ANSWER: c 69. How does the formation of antibonding molecular orbitals affect the stability of a molecule? a. It increases the stability of the molecule. b. It decreases the stability of the molecule. c. It has no effect on the stability of the molecule. d. It decreases the reactivity of the molecule. e. It stabilizes the molecule in specific conditions.

ANSWER: b 70. Explain how certain molecules exhibit paramagnetism according to molecular orbital theory.

ANSWER Paramagnetism occurs when molecules have unpaired electrons in their : molecular orbitals. This property arises from the presence of magnetic moments due to unpaired spins.

71. How do bonding and antibonding molecular orbitals form according to molecular orbital (MO) theory?

ANSWER In MO theory, atomic orbitals merge to create molecular orbitals, which are : mathematical descriptions of electron waves across a molecule. Bonding orbitals (with lower energy) and antibonding orbitals (with higher energy) can result from the combination of atomic orbitals.

72. What is the significance of the bond order in molecular orbital theory and its relationship to bond strength and stability?

ANSWER Bond order reflects the number of electrons in bonding versus antibonding : orbitals. Higher bond orders indicate stronger bonds and greater stability.

73. Describe how the number of nodes in a molecular orbital affects its energy and stability.

ANSWER Nodes represent regions where electron probability density is zero. The more : nodes in a molecular orbital, the higher its energy and the less stable it is.

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Chapter 9 Molecular Shape, Polarity, and Advanced Bonding Models 74. Draw the MO diagram for H2 and calculate the bond order for the molecule.

ANSWER:

The bond order is 1.

75. Explain how the concept of constructive and destructive interference impacts bonding in molecular orbital theory.

ANSWER Constructive interference leads to the formation of bonding molecular orbitals : with lower energy, while destructive interference forms antibonding molecular orbitals with higher energy.

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Chapter 8 Molecular Structure and Bond Energy 1. In the Lewis structure for ammonia (NH3), how many lone pairs of electrons are present on the nitrogen atom? a. 0 b. 1 c. 2 d. 3 e. 4 ANSWER: b 2. Which of the following molecules contains a double bond? a. H2O b. CH4 c. N2 d. CO2 e. HF ANSWER: d 3. Which of the following molecules does NOT contain a double or triple covalent bond? a. HCN b. Cl2 c. N2 d. CO2 e. O2 ANSWER: b 4. Sulfur has 12 valence electrons in SF6. Which of the following is TRUE about sulfur in this molecule? a. Sulfur has a formal charge of –1. b. Sulfur has a formal charge of +1. c. Sulfur has an expanded octet. d. Sulfur has one double bond and five single bonds. e. Sulfur has two double bonds and four single bonds.

ANSWER: c 5. Which of the following elements can be represented by the Lewis dot symbol shown?

a. silicon Copyright Macmillan Learning. Powered by Cognero.

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Chapter 8 Molecular Structure and Bond Energy b. sulfur c. boron d. iodine e. argon ANSWER: d 6. Which of the following elements can be represented by the Lewis dot symbol shown?

a. phosphorous b. silicon c. aluminum d. chlorine e. oxygen ANSWER: e 7. Which of the following elements can be represented by the Lewis dot symbol shown?

a. nitrogen b. carbon c. magnesiu m d. fluorine e. aluminum ANSWER: b 8. What is the formal charge of oxygen in the hydronium ion?

a. 0 b. + 2 c. + 1 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 8 Molecular Structure and Bond Energy d. –1 e. –2 ANSWER: c 9. What is the formal charge of carbon in carbon monoxide?

a. 0 b. + 2 c. + 1 d. –1 e. –2 ANSWER: d 10. What is the formal charge of oxygen in carbon monoxide?

a. 0 b. + 2 c. + 1 d. –1 e. –2 ANSWER: c 11. What is the octet rule?

ANSWER The octet rule states that atoms with eight electrons in their highest occupied : energy (valence) level are very stable.

12. Draw the Lewis dot symbol for aluminum.

ANSWER:

13. Draw the Lewis dot symbol for fluorine.

ANSWER:

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Chapter 8 Molecular Structure and Bond Energy 14. Draw the Lewis dot symbol for nitrogen.

ANSWER:

15. Calculate the formula charge for each atom in the nitrite ion.

ANSWER: left O: –1 N: 0 right O: 0 16. How many valence electrons are in CCl4? a. 30 b. 32 c. 18 d. 26 e. 8 ANSWER: b 17. How many valence electrons are in NH4+? a. 5 b. 6 c. 7 d. 8 e. 9 ANSWER: d 18. Which of the following statements describes the bonding arrangement of phosphorous in the molecule PH3? a. Phosphorus has three single bonds to the H atoms. b. Phosphorus has three single bonds to the H atoms and one lone pair of electrons. c. Phosphorus has two single bonds, one double bond to the H atoms, and one lone pair of electrons. d. Phosphorus has two single bonds and one double bond to the H atoms. e. Phosphorus has one single bond and two double bonds to the H atoms.

ANSWER: b

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Chapter 8 Molecular Structure and Bond Energy 19. Which of the following statements describes the bonding arrangement of sulfur in the molecule H2S? a. Sulfur has two single bonds to the H atoms and two lone pairs of electrons. b.Sulfur has two double bonds to the H atoms. c. Sulfur has two single bonds to the H atoms and one lone pair of electrons. d.Sulfur has a single bond to an H atom, a double bond to the second H atom, and one lone pair of electrons. e. Sulfur has two double bonds to the H atoms and two lone pairs of electrons.

ANSWER: a 20. Which of the following statements describes the bonding arrangement of carbon in HCO2–? a. Carbon has two single bonds and one triple bond. b. Carbon has one single bond and two double bonds. c. Carbon has three single bonds and one lone pair of electrons. d. Carbon has three single bonds and two lone pairs of electrons. e. Carbon has two single bonds and one double bond.

ANSWER: e 21. Which of the following statements describes the bonding arrangement of phosphorous in PH4+? a. Phosphorous has one double bond and three single bonds. b. Phosphorous has two double bonds and two single bonds. c. Phosphorous has three double bonds and one single bond. d. Phosphorous has four single bonds. e. Phosphorous has four single bonds and one lone pair of electrons.

ANSWER: d 22. Which of the following Lewis structures represents CO32–? a.

b.

c.

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Chapter 8 Molecular Structure and Bond Energy d.

e.

ANSWER: d 23. Which of the following Lewis structures represents methanol, CH3OH? a.

b.

c.

d.

e.

ANSWER: b 24. What is the chemical formula of a molecule formed when a single selenium atom reacts with hydrogen gas?

Assume that the selenium atom follows the octet rule. a. HSe b. H2Se c. SeH3 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 8 Molecular Structure and Bond Energy d. SeH4 e. SeH5 ANSWER: b 25. What is the chemical formula of a molecule formed when a single silicon atom reacts with hydrogen gas?

Assume that the silicon atom follows the octet rule. a. SiH b. SiH2 c. SiH3 d. SiH4 e. SiH5 ANSWER: d 26. What is the chemical formula of a molecule formed when a single arsenic atom reacts with hydrogen gas?

Assume that the arsenic atom follows the octet rule. a. AsH b. AsH2 c. AsH3 d. AsH4 e. AsH5 ANSWER: c 27. Molecules that have the same molecular formula but different structural formulas are known as a. isomers. b. isotopes. c. allotropes. d. isobars. e. isoelectronic.

ANSWER: a 28. Which of the following Lewis structures represents the most stable form of sulfate, SO42–? a.

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Chapter 8 Molecular Structure and Bond Energy b.

c.

d.

e.

ANSWER: d 29. Which of the following Lewis structures represents acetic acid, HC2H3O2? a.

b.

c.

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Chapter 8 Molecular Structure and Bond Energy d.

e.

ANSWER: b 30. What is the normal bonding arrangement of oxygen with a formal charge of zero? a. 1 bond, 0 lone pairs b. 4 bonds, 0 lone pairs c. 3 bonds, 1 lone pair d. 2 bonds, 2 lone pairs e. 1 bond, 3 lone pairs ANSWER: d 31. What is the normal bonding arrangement of chlorine with a formal charge of zero? a. 1 bond, 0 lone pairs b. 4 bonds, 0 lone pairs c. 3 bonds, 1 lone pair d. 2 bonds, 2 lone pairs e. 1 bond, 3 lone pairs ANSWER: e 32. Draw the Lewis structure for ethanol, CH3CH2OH. ANSWER:

33. Draw the Lewis structure for hydrazine, N2H4.

ANSWER:

34. Draw the Lewis structure for dichloromethane, CH2Cl2.

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Chapter 8 Molecular Structure and Bond Energy ANSWER:

35. Draw the Lewis structure for sulfur dioxide, SO2.

ANSWER:

36. Draw the Lewis structure for hypochlorite, ClO–. Determine the formal charge for each atom.

ANSWER:

Formal charges: O = –1 Cl = 0 37. Draw the Lewis structure for phosphate, PO43–, that minimizes the formal charge across the ion.

ANSWER:

38. Draw the Lewis structure for nitric acid, HNO3.

ANSWER:

39. Draw the two constitutional isomers for a molecule with the formula C3H7F. ANSWER:

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Chapter 8 Molecular Structure and Bond Energy 40. What is the average formal charge on each oxygen atom in sulfite, SO32–? a. –2 b. –1 c. –1/3 d. –2/3 e. –1/2 ANSWER: d 41. How many resonance structures are possible for the perchlorate ion, ClO4–? a. 0 b. 1 c. 2 d. 3 e. 4 ANSWER: e 42. Draw the three resonance structures for the thiocyanate ion, SCN–.

ANSWER :

43. What is the condensed formula for octane, CH3CH2CH2CH2CH2CH2CH2CH3? a. CH3(CH2)4CH3 b. CH3(CH2)5CH3 c. CH3(CH2)6CH3 d. CH3(CH2)7CH3 e. CH3(CH2)8CH3 ANSWER: c 44. What is the molecular formula of a molecule with a condensed formula of CH3(CH2)5CH3? a. C7H18 b. C7H16 c. C7H14 d. C6H16 e. C6H14 ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 8 Molecular Structure and Bond Energy 45. What is the condensed formula for 2-aminopentane?

a. CH3CHNH2(CH2)2CH3 b. CH3NH2(CH2)3CH3 c. CH3(CH)2NH2(CH2)2CH3 d. CH3CH2NH2(CH2)3CH3 e. CH3NH2(CH2)4CH3 ANSWER: a 46. What is the condensed formula of butane?

a. CH3CH2CH2CH2CH3 b. CH3(CH2)4CH3 c. CH3(CH2)3CH3 d. CH3(CH2)2CH3 e. (CH3)2(CH2)2 ANSWER: d 47. What is the molecular formula of the molecule whose skeletal formula is shown?

a. C3H6 b. C3H8 c. C3H10 d. C4H10 e. C4H12 ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 8 Molecular Structure and Bond Energy 48. What is the condensed formula of the molecule whose skeletal formula is shown?

a. CH3(CH2)6CH3 b. CH3(CH2)5CH3 c. CH3(CH2)4CH3 d. CH3(CH2)3CH3 e. CH3(CH2)8CH3 ANSWER: c 49. What is the molecular formula of the molecule whose skeletal formula is shown?

a. C6H7O b. C6H12O c. C7H7O d. C6H6O e. C6H8O ANSWER: d 50. Draw the corresponding Lewis structure for the following molecule.

ANSWER:

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Chapter 8 Molecular Structure and Bond Energy 51. Draw the corresponding skeletal structure for the following molecule.

ANSWER:

52. Draw the skeletal structure for the following molecule with the condensed formula of CH3(CH2)3CH3.

ANSWER:

53. Estimate the enthalpy change (ΔH) for the following reaction using the bond energy data provided.

H2(g) + Cl2(g) → 2HCl(g) Bond Energies (kJ/mol): H–H = 436 Cl–Cl = 243 H–Cl = 432 a. –185 kJ b. 185 kJ c. –247 kJ d. 247 kJ e. –239 kJ ANSWER: a 54. What is the enthalpy change for the combustion of methane (CH4) based on the given bond energies?

CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) Bond Energies (kJ/mol): C–H = 410 O=O = 498 C=O = 745 O–H = 460 a. –226 kJ b. 226 kJ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 8 Molecular Structure and Bond Energy c. –694 kJ d. 694 kJ e. –532 kJ ANSWER: c 55. Based on the bond energies, which of the following bonds is the strongest?

Bond Energy (kJ/mol) H–O 460 H–F 569 H–Cl 432 H–Br 366 H–I 298 a. H–O b. H–F c. H–Cl d. H–Br e. H–I ANSWER: b 56. Based on the bond energies, which of the following bonds is the weakest?

Bond Energy (kJ/mol) F–F 157 Cl–Cl 243 Br–Br 194 I–I 153 H–H 436 a. F–F b. Cl–Cl c. Br–Br d. I–I e. H-H ANSWER: d 57. Which of following statements is INCORRECT? a. It always requires energy to break a chemical bond. b.Energy is always released when chemical bonds form. c. The ΔH for a reaction is the sum of the enthalpy changes for the bonds broken and the bonds formed.

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Chapter 8 Molecular Structure and Bond Energy d.The bond energy associated with a particular bond will be the same across all molecules that contain that bond. e. Bond energies for bond broken in reaction have positive values.

ANSWER: d 58. What types of bonds are broken in the combustion of formaldehyde, CH2O, in which it reacts with oxygen in the atmosphere?

ANSWER: 1 C=O, 2 C–H, and 1 O=O bond 59. List all of the bonds that are broken and formed in the following chemical reaction.

CH2Cl2(g) + O2(g) → CO2(g) + 2 HCl(g) ANSWER: Bonds broken: 2 C–H, 2 C–Cl, and 1 O=O bond Bonds formed: 2 C=O and 2 H–Cl bonds 60. Estimate the enthalpy change for the decomposition of hydrogen peroxide given the following bond energies.

2H2O2(g) → 2H2O(g) + O2(g) Bond Energies (kJ/mol): H–O 460 H–H 436 O=O 498 O–O 142 ANSWER: –214 kJ 61. Which of the following statements correctly defines lattice energy? a. The energy required to break apart an ionic solid into its constituent gaseous ions. b. The energy released when an ionic solid forms from its constituent gaseous ions. c. The energy required to convert a mole of ionic solid into its constituent ions in solution. d. The energy required to remove an electron from a gaseous atom. e. The energy required to add an electron to a gaseous atom.

ANSWER: a 62. Which of the following ionic compounds is expected to have the highest lattice energy? a. LiF b. NaCl c. KBr d. MgO Copyright Macmillan Learning. Powered by Cognero.

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Chapter 8 Molecular Structure and Bond Energy e. CaS ANSWER: d 63. Which of the following ionic compounds is expected to have the lowest lattice energy? a. MgO b. CaO c. BaO d. SrO e. BeO ANSWER: c 64. Which of the following processes is not involved in the Born-Haber cycle for the formation of lithium fluoride (LiF)? a. ionization of lithium b. ionization of fluorine c. sublimation of lithium d. dissociation of fluorine e. hydration of lithium fluoride ANSWER: e 65. Which of the following ionic compounds is expected to have the highest melting point? a. LiCl b. NaCl c. KCl d. RbCl e. CsCl ANSWER: a 66. Which of the following ionic compounds is expected to have the highest solubility in water? a. LiF b. NaF c. KF d. RbF e. CsF ANSWER: e 67. Which of the following statements best explains why the lattice energy of MgO is greater than that of NaF? a. Mg2+ has a larger ionic radius than Na+. b. O2– has a smaller ionic radius than F–. c. Mg2+ and O2– have higher charges than Na+ and F–.

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Chapter 8 Molecular Structure and Bond Energy d. MgO has a crystal structure that is different from NaF. e. MgO is more covalent than NaF. ANSWER: c 68. Which of the following statements correctly explains why the lattice energy of NaCl is greater than that of KCl? a. Na+ has a smaller ionic radius than K+. b. Cl– has a larger ionic radius than K+. c. NaCl have higher charges than KCl. d. NaCl has a crystal structure that is different from KCl. e. NaCl is more covalent than KCl. ANSWER: a 69. Rank the following ionic compounds in order of increasing lattice energy: NaCl, CaF2, LiF, CaCl2. a. NaCl < CaF2 < LiF < CaCl2 b. NaCl < LiF < CaF2 < CaCl2 c. NaCl < LiF < CaCl2 < CaF2 d. CaCl2 < CaF2 < NaCl < LiF e. CaF2 < CaCl2 < LiF < NaCl ANSWER: c 70. Rank the following ionic compounds in order of decreasing melting point: KCl, MgF2, RbF, MgCl2. a. KCl > MgF2 > RbF > MgCl2 b. KCl > RbF > MgF2 > MgCl2 c. KCl > RbF > MgCl2 > MgF2 d. MgCl2 > MgF2 > KCl > RbF e. MgF2 > MgCl2 > RbF > KCl ANSWER: e 71. What are the three stages of the Born–Haber cycle?

ANSWER The three stages are the formation of gaseous atoms, the formation of gaseous : ions, and the assembly of the ions to form an ionic lattice.

72. What processes of the Born–Haber cycle require energy? What processes release energy?

ANSWER Processes that require energy: : 1. Converting a solid metal to gaseous atoms.

2. Bond dissociation of the nonmetal molecule.

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Chapter 8 Molecular Structure and Bond Energy 3. Removing one or more electrons from the metal atom to form a cation.

Processes that release energy: 1. Adding one or more electrons to the nonmetal atom to form an anion.

2. Energy is released when the cations and anions combine to form the ionic solid.

73. Using the following information, calculate the lattice energy for zinc fluoride, ZnF2.

Ionization Energy (kJ/mol) IE1 IE2 Zn 906 1,733 Electron Affinity (kJ/mol) F –328 Standard Enthalpies of Formation (kJ/mol) Zn(s) 0 Zn(g) 130.4 –764 ZnF2(s) 0 F2(g) F(g) 79.4 ANSWER: −3,036 kJ/mol 74. Using the following information, calculate the lattice energy for magnesium iodide, MgI2.

Ionization Energy (kJ/mol) IE1 IE2 Mg 738 1,450 Electron Affinity (kJ/mol) I –295 Standard Enthalpies of Formation (kJ/mol) Mg(s) 0 Mg(g) 147.1 –364 MgI2(s) 0 I2(g) I(g) 106.8 ANSWER: −2,323 kJ/mol 75. What factors affect the lattice energy of ionic compounds? Explain the relationship between the factors and the lattice energy.

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Chapter 8 Molecular Structure and Bond Energy ANSWER The two factors that affect lattice energy are the ion’s charge and radius. The : greater the charge of the ions, the stronger the electrostatic attraction between them, leading to a higher lattice energy. Smaller ions result in stronger electrostatic attractions and, therefore, higher lattice energy.

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Chapter 7 Quantum Mechanics and Electronic Structure 1. What measurement describes the number of waves that pass a point over a period of time? a. wavelength b. speed c. frequency d. energy e. photoelectric effect ANSWER: c 2. What measurement describes the distance between the same points on adjacent waves? a. wavelength b. speed c. frequency d. energy e. photoelectric effect ANSWER: a 3. What is the frequency of a green laser pointer that has a wavelength of 532 nm? a. 1.97×1014 Hz b. 5.32×1014 Hz c. 6.19×1014 Hz d. 5.64×1014 Hz e. 7.32×1014 Hz ANSWER: d 4. What is the wavelength of the radio waves emitted for a station with a frequency of 97.9 MHz? a. 3.06×107 m b. 1.21×107 m c. 6.15 m d. 1.21 m e. 3.06 m ANSWER: e 5. What is the energy of an X-ray photon that has a wavelength of 0.25 nm? a. 4.6×10–16 J b. 2.5×10–17 J c. 1.7×10–19 J d. 4.2×10–17 J Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure e. 8.0×10–16 J ANSWER: e 6. Which of the following statements is TRUE? a. The speed of light is proportional to energy. b. Energy is inversely proportional to frequency. c. Wavelength is inversely proportional to frequency. d. Energy is proportional to wavelength. e. The energy of electromagnetic radiation is constant.

ANSWER: c 7. What is the energy of a UV photon that has a wavelength of 275 nm? a. 5.60×10–19 J b. 7.23×10–19 J c. 8.21×10–19 J d. 4.07×10–19 J e. 1.69×10–19 J ANSWER: b 8. A laser emits a pulse of light with an energy of 30.5 mJ at a wavelength of 1064 nm. How many photons are emitted in the pulse? a. 1.63×1017 photons b. 4.05×1019 photons c. 9.17×1019 photons d. 2.36×1019 photons e. 1.23×1020 photons ANSWER: a 9. The minimum energy required to remove an electron from tin is 7.08 × 10–19 J. Which of the following frequencies would be able to remove an electron from tin? a. 48.2 MHz b. 1070 THz c. 984 GHz d. 498 GHz e. none of the above ANSWER: b 10. Which electron energy level transitions would emit a photon with the highest frequency?

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Chapter 7 Quantum Mechanics and Electronic Structure a. 2 → 1 b. 3 → 1 c. 4 → 1 d. 1 → 5 e. 1 → 6 ANSWER: c 11. What is the photoelectric effect?

ANSWER The photoelectric effect is the phenomenon in which electrons are emitted from : a material's surface when it absorbs light of sufficient energy.

12. A broad-spectrum sunscreen provides protection from ultraviolet (UV) light with wavelengths ranging from 100 to 400 nanometers (nm). Determine the range of the frequency and energy of this light.

ANSWER frequency range: 3.00×1015 to 7.50×1014 Hz : energy range: 4.97×10−19 to 1.99×10−18 J 13. A flame test was conducted on a solution containing an unknown metal ion from the following list. Light with a frequency of 5.43×1014 Hz was observed. Using the following information, what is the wavelength (in nm) of light observed and the identity of the metal ion present in the solution?

Metal Sodium Potassium Barium Strontium

Flame Test Color yellow violet green red

Color Violet Blue Green Yellow Orange Red

Wavelengths (nm) 380–450 450–495 495–570 570–590 590–620 620–750

ANSWER: λ = 552 nm The identity of the metal ion is barium.

14. Argon has an ionization energy of 1450 kJ/mol. What is the minimum frequency of light that can ionize a single argon atom?

ANSWER: 3.64×1015 Hz Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure 15. What is a flame test? How can it be used to identify an element?

ANSWER A flame test is a procedure used to detect the presence of metal ions based on : the characteristic color they emit when heated in a flame. During a flame test, a solution containing a metal ion is placed in a flame, and the electrons in the metal ions absorb energy, becoming excited to higher energy levels. When these electrons return to their ground state, they release energy in the form of light, with a color characteristic of the metal ion.

16. What is the wavelength of an electron traveling at 2.5×106 m/s? The mass of an electron is 9.11×10–31 kg. a. 3.6×10–10 m b. 4.2×10–10 m c. 1.9×10–9 m d. 2.9×10–10 m e. 4.0×10–9 m ANSWER: d 17. What is the wavelength of a tennis ball served at 195 km/h? The mass of the tennis ball is 56.0 g. a. 1.9×10–38 m b. 4.9×10–38 m c. 6.1×10–38 m d. 2.2×10–34 m e. 1.3×10–39 m ANSWER: c 18. Which particle would have the shortest wavelength when traveling at 10% of the speed of light? a. electron, 9.1094×10–31 kg b. proton, 1.6726×10–27 kg c. neutron, 1.6749×10–27 kg d. muon, 1.8835×10–28 kg e. tau, 3.1675×10–27 kg ANSWER: e 19. Which of the following phenomena is most directly explained by the wave nature of electrons? a. the diffraction of electrons by a crystal b. the quantization of energy levels in an atom c. the photoelectric effect d. the emission of photons from excited atoms Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure e. the radioactive decay of unstable nuclei ANSWER: a 20. Which of the following is a consequence of Heisenberg's uncertainty principle? a. The energy of an electron is quantized. b. The quantization of energy levels in an atom. c. Electrons exhibit wave-particle duality. d. The exact position and momentum of an electron cannot be known simultaneously. e. Electrons in an atom occupy specific energy levels.

ANSWER: d 21. Which of the following statements about wave-particle duality is INCORRECT? a. All matter exhibits wave-particle duality. b. The wave nature of matter is more pronounced for macroscopic objects. c. The de Broglie wavelength of an object is inversely proportional to its speed. d. The wave-particle duality is a fundamental concept in quantum mechanics. e. The double-slit experiment can be used to demonstrate the wave nature of electrons.

ANSWER: b 22. The diameter of a silicon atom used in a semiconductor is 210 pm. If the position of an electron is known within this diameter, determine the minimum uncertainty in its speed. The mass of the electron is 9.11×10–31 kg. a. 7.0×105 m/s b. 9.5×105 m/s c. 2.8×105 m/s d. 1.3×106 m/s e. 5.2×106 m/s ANSWER: c 23. The diameter of a hydrogen atom is approximately 120 pm. If the position of an electron is known within this diameter, what is the minimum uncertainty in the electron's speed? The mass of the electron is 9.11×10–31 kg.

ANSWER: 4.8×105 m/s 24. When a radioactive carbon-14 nucleus decays, it emits a particle with a mass of 9.11×10–31 kg and a speed of 2.15×107 m/s. What is the de Broglie wavelength of this particle? If the uncertainty in the momentum (mΔu) is 2.1×10−21 kg·m/s, what is the minimum possible uncertainty in the position of this particle?

ANSWER: λparticle = 3.38×10–11 m Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure Δx = 2.5×10–14 m 25. What occurs to waves when they undergo constructive and destructive interference?

ANSWER Constructive interference occurs when waves meet in phase, meaning their : crests and troughs align. When this happens, the amplitudes of the waves add together, resulting in a wave with a larger amplitude than either of the original waves. Destructive interference occurs when waves meet out of phase, meaning the crest of one wave aligns with the trough of another. In this case, the amplitude of the resulting wave is reduced or canceled out.

26. How many electrons can occupy a p sublevel? a. 2 b. 3 c. 6 d. 10 e. 14 ANSWER: c 27. How many orbitals are there in an f sublevel? a. 1 b. 3 c. 6 d. 7 e. 14 ANSWER: d 28. What is the maximum number of electrons that can be present within the third energy level? a. 32 b. 18 c. 14 d. 8 e. 2 ANSWER: b 29. Which of the following values is NOT a possible value for the principal quantum number? a. 0 b. 1 c. 2 d. 5 e. 10 ANSWER: a Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure 30. Which of the following values is a solution for the spin quantum number? a. 0 b. 1 c. –2 d. –1 e. –½ ANSWER: e 31. Which of the following sets of quantum numbers is possible for an electron in a 3p orbital? a. [3, 2, –1, +½] b. [4, 3, –2, +½] c. [3, 2, –1, +½] d. [3, 0, 0, –½] e. [3, 1, 0, +½] ANSWER: e 32. What are the possible ml values for an electron in a 3d orbital? a. –1, 0, 1 b. –2, –1, 0, 1, 2 c. –3, –2, –1, 0, 1, 2, 3 d. +½, –½ e. 3, 2, 1, 0 ANSWER: b 33. Which of the following sets of quantum numbers is NOT possible? a. [5, 2, –1, –½] b. [4, 0, 0, +½] c. [3, 2, –1, +½] d. [3, 0, –1, –½] e. [4, 2, 0, +½] ANSWER: d 34. What is the Pauli exclusion principle?

ANSWER The Pauli exclusion principle states that no two electrons in a single atom can : have the same four quantum numbers.

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Chapter 7 Quantum Mechanics and Electronic Structure 35. What are the possible quantum numbers for an electron in a 4s orbital?

ANSWER: [4, 0, 0, +½] or [4, 0, 0, –½] 36. What is the level and sublevel of an electron with the quantum numbers of [3, 2, 0, –½]?

ANSWER: 3 d 37. How many orbitals are possible that have a principal quantum number of 4? How many electrons can occupy those orbitals?

ANSWER: There are 16 possible orbitals in n = 4. 32 electrons can occupy the 16 orbitals.

38. Which element has the following electron configuration as a neutral atom?

1s22s22p4 a. C b. N c. O d. F e. S ANSWER: c 39. Which element has the following electron configuration as a neutral atom?

1s22s22p6 a. C b. N c. O d. F e. Ne ANSWER: e 40. What is the electron configuration for a strontium atom? a. [Ar]4s23d104p6 b. [Ar]5s24d105p6 c. [Kr]5s2 d. [Kr]5s24d105p6 e. [Ar]5s24d1 ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure 41. Which of the following electron configurations represents an atom with 7 valence electrons? a. [Ar]4s23d104p6 b. [Ar]5s24d105p5 c. [Kr]5s2 4d7 d. [Kr]5s24d105p3 e. [Ar]5s24d5 ANSWER: b 42. How many valence electrons are present in an arsenic atom? a. 2 b. 5 c. 8 d. 3 e. 15 ANSWER: b 43. How many valence electrons are present in a tin atom? a. 2 b. 5 c. 8 d. 4 e. 14 ANSWER: d 44. What is the electron configuration for a phosphide ion, P3–? a. 1s22s22p63s23p3 b. 1s22s22p63s23p4 c. 1s22s22p63s23p5 d. 1s22s22p63s23p6 e. 1s22s22p63s23p64s23d1 ANSWER: d 45. Which of the following ions is NOT isoelectronic with argon? a. P3– b. S2– c. Cl– d. K+ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure e. Na+ ANSWER: e 46. What is the electron configuration for a calcium ion, Ca2+?

ANSWER: 1s22s22p63s23p6 47. What is the electron configuration for a neutral nitrogen atom?

ANSWER: 1s22s22p3 48. Write the electron configuration for bromine using noble gas shorthand. How many inner, outer, and valence electrons does bromine have?

ANSWER: [Ar]4s23d104p5 Bromine has 18 inner electrons, 17 outer electrons, and 7 valence electrons.

49. Write the electron configuration for titanium using noble gas shorthand. How many inner, outer, and valence electrons does titanium have?

ANSWER: [Ar]4s23d2 Titanium has 18 inner electrons, 4 outer electrons, and 2 valence electrons.

50. How many valence electrons does carbon have? a. 2 b. 4 c. 6 d. 8 e. 0 ANSWER: b 51. How many valence electrons do elements in group 17 (7A) have? a. 2 b. 4 c. 6 d. 7 e. 8 ANSWER: d 52. How many valence electrons do elements in group 2 (2A) have? a. 2 b. 4 c. 6 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure d. 7 e. 8 ANSWER: a 53. Which of the following elements have a valence-level electron configuration of s2p1? a. Be b. Na c. K d. Sc e. Al ANSWER: e 54. Which of the following elements have a valence-level electron configuration of s2p2? a. B b. S c. Ti d. Si e. Al ANSWER: d 55. What is the electron configuration of Zn2+? a. [Ar]4s23d10 b. [Ar]3d10 c. [Ar]4s23d8 d. [Ar]4s23d104p2 e. [Ar]4s13d9 ANSWER: b 56. What is the electron configuration for the highest-energy occupied sublevel of antimony? a. 3d10 b. 4p4 c. 4p3 d. 4s2 e. 3d9 ANSWER: c 57. Which group of elements has 6 valence electrons?

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Chapter 7 Quantum Mechanics and Electronic Structure a. group 2 (2A) b. group 4 (4A) c. group15 (5A) d. group 16 (6A) e. group 18 (8A) ANSWER: d 58. Based on the periodic table, molybdenum’s electron configuration is expected to be [Kr]5s24d4, but it is actually [Kr]5s14d5. What causes this to happen?

ANSWER Half-filled s and d sublevels are slightly more stable than a full s orbital and : partially filled d sublevel.

59. Based on the periodic table, silver's electron configuration is expected to be [Kr]5s24d9, but it is actually [Kr]5s14d10. What causes this to happen?

ANSWER A half-filled s and a full d sublevel are doubly stabilized and more favorable : than a full s orbital and a partially filled d sublevel.

60. Using the periodic table, chromium's electron configuration is expected to be [Ar]4s23d4, but it has been experimentally determined to have a different electron configuration. What is chromium's actual electron configuration? Why does it deviate from the expected electron configuration?

ANSWER Chromium's actual electron configuration is [Ar]4s13d5. The half-filled s and d : sublevels are slightly more stable than a full s orbital with a partially filled d sublevel.

61. Using the periodic table, copper's electron configuration is expected to be [Ar]4s23d9, whereas it has been experimentally determined to have a different electron configuration. What is copper's actual electron configuration? Why does it deviate from the expected electron configuration?

ANSWER Copper's actual electron configuration is [Ar]4s13d10. A half-filled s and a : full d sublevel are doubly stabilized and more favorable than a full s orbital and a partially filled d sublevel.

62. What are the electron configurations for Ge, Ge2+, and Ge4+?

ANSWER: Ge: [Ar]4s23d104p2 Ge2+: [Ar]4s23d10 Ge4+: [Ar]3d10 63. What is the effective nuclear charge for a fluorine atom? a. + 9 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure b. + 8 c. + 7 d. + 6 e. + 2 ANSWER: c 64. Which element has the highest effective nuclear charge? a. Mg b. S c. B d. C e. N ANSWER: b 65. Which of the following elements has the largest atomic radius? a. B b. Bi c. Ba d. Br e. Be ANSWER: c 66. Which of the following elements has the smallest atomic radius? a. Ne b. S c. Li d. As e. Sn ANSWER: a 67. Which of the following isoelectronic ions has the largest radius? a. Cl– b. S2– c. P3– d. K+ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure e. Ca2+ ANSWER: c 68. Which element has the highest ionization energy? a. Fe b. S c. Al d. O e. N ANSWER: d 69. Which of the following equations represents the second ionization energy of oxygen? a. O(g) + 2e– → O2–(g) b. O–(g) + e– → O2–(g) c. O2–(g) + 2e–→ O(g) d. O(g) → O2+(g) + 2e– e. O+(g) → O2+(g) + e– ANSWER: e 70. What property is a measure of the energy change when an atom gains an electron? a. ionization energy b. effective nuclear charge c. metallic character d. electron affinity e. nonmetallic character ANSWER: d 71. The electron affinity of chlorine is –349 kJ/mol. What is the change in energy associated with 10.0 g of Cl reacting to produce Cl–? a. –109 kJ b. –98.4 kJ c. –56.9 kJ d. –117 kJ e. –72.3 kJ ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 7 Quantum Mechanics and Electronic Structure 72. Write the equations that represent the first and second ionization energies of sulfur.

ANSWER: first ionization energy: S(g) → S+(g) + e– second ionization energy: S+(g) → S2+(g) + e– 73. Rank the following elements in order of increasing atomic radius: Se, Ca, N, O, K. ANSWER: O < N < Se < Ca < K 74. Rank the following isoelectronic species in order of decreasing radius: Xe, Cs+, Ba2+, I–, Te2–.

ANSWER: Te2– > I– > Xe > Cs+ > Ba2+ 75. What is the metallic character of an element? Describe the periodic trends for metallic character for periods and groups.

ANSWER The metallic character of an element refers to whether or not its chemical and : physical properties are associated with metals. Elements with high metallic character are typically good conductors of heat and electricity, shiny, malleable, and tend to lose electrons easily in chemical reactions. As you move across a period from right to left on the periodic table, the metallic character increases. As you move down a group, the metallic character also increases.

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Chapter 6 Energy 1. Which is the official IUPAC unit of energy? a. joules (J) b. calories (cal) c. Calories (Cal) d. British Thermal Units (BTU) e. kilowatt-hour (kWh) ANSWER: a 2. Which equation correctly relates heat (q), work (w) and the energy change (ΔE)? a. b. c. d. e.

ANSWER: d 3. A sample of a flammable gas is burned, and it gives off 200 J as heat and does 50 J of work. What is the total energy change for the system? a. +250 J b. –250 J c. –150 J d. +150 J e. 0 J ANSWER: c 4. The work (w) done by an expanding gas is a. positive. b. negative. c. zero. d. dependent on ΔE. e. dependent on q.

ANSWER: b 5. What is heat (q) equal to at constant pressure? a. work b. PV work c. internal energy d. enthalpy e. temperature Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy ANSWER: d 6. A piston expands against a constant pressure of 3.1 atm from a volume of 1.2 L to 5.4 L. What is the work done by the surroundings for this process? a. –1.3 kJ b. +1.3 kJ c. –2.6 kJ d. +2.6 kJ e. +3.1 kJ ANSWER: b 7. What is the equation for pressure-volume work? a. b. c. d. e.

ANSWER: c 8. Which of the following best describes the first law of thermodynamics? a. Energy can only be created, not destroyed. b. Energy is equal to work plus heat. c. Internal energy is always conserved. d. The total energy of the universe is constant. e. Energy cannot be created or destroyed but can change forms.

ANSWER: e 9. What does the first law of thermodynamics simplify to if ΔE = 0 for a process? a. q = w b. q = 0 c. w = 0 d. q = ΔH e. q + w = ΔH ANSWER: a 10. A bag of potato chips contains 160 kilocalories of stored energy. What is this in kilojoules? a. 0.67 kJ b. 38 kJ c. 140 kJ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy d. 670 kJ e. 160 kJ ANSWER: d 11. What is the definition of pressure? a. pressure = energy/area b. pressure = energy/force c. pressure = force/area d. pressure = force/volume e. pressure = force×area ANSWER: c 12. A chemical reaction produces 34.1 J of energy. What is this in calories? a. 0.800 cal b. 6.70 cal c. 8.15 cal d. 14.3 cal e. 143 cal ANSWER: c 13. When a sample of methane is combusted, it releases 34.5 kJ of energy. If the sample does 14.2 kJ of work, how much energy is lost as heat? a. 20.3 kJ b. –20.3 kJ c. 48.7 kJ d. –48.7 kJ e. –34.5 kJ ANSWER: b 14. What is the sign of the change in internal energy of the surroundings when the change in internal energy of the system is positive?

ANSWER: negative 15. Calculate the work done by a gas that expands from 2.00 L to 6.00 L against a constant pressure of 1.50 atm.

ANSWER: –608 J 16. Water condenses on a pane of glass. Is this process exothermic or endothermic? Explain your answer.

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Chapter 6 Energy ANSWER: Heat is released when water condenses, so this is an exothermic process.

17. A candle performs 83.7 J of work, while releasing 8.90 J of heat. What is the change in internal energy for this system?

ANSWER: –92.6 J 18. For a certain process, the change in the internal energy of the system is –3.2 kJ. What is the change in the internal energy of the surroundings?

ANSWER: +3.2 kJ 19. In a certain physical process, the internal energy of the system changes by +5.4 kJ. If the process also performs 3.3 kJ of work on its surroundings, how much heat is produced or absorbed? Is the reaction exothermic or endothermic?

ANSWER: +8.7 kJ; endothermic 20. A chemical reaction releases 155 J of heat while also expanding a cylinder by 1.33 L against 1.01 atm of pressure. What is the change of internal energy for the system?

ANSWER: ∆E = –291 J 21. An exothermic reaction releases 1327 kJ as heat. How many Calories (kcal) is this?

ANSWER: 317.2 Cal 22. What is the difference between heat and temperature? a. Heat measures molecular motion, whereas temperature measures energy transfer. b. Heat is energy in transit, whereas temperature is a measure of average kinetic energy. c. Heat depends on mass, whereas temperature is constant for all objects. d. Heat is stored energy, whereas temperature measures potential energy. e. Heat flows in cycles, whereas temperature remains fixed.

ANSWER: b 23. How does heat transfer occur without additional work input? a. From an object with higher mass to one with lower mass. b. From a colder object to a warmer object. c. From a warmer object to a cooler object. d. From objects of higher density to lower density. e. Only in the presence of a medium.

ANSWER: c Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy 24. How much heat does it take to increase the temperature of 50.0 g of water from 22.0 ºC to 50.0 ºC? a. 2210 J b. 3930 J c. 4600 J d. 5850 J e. 10500 J ANSWER: d 25. Suppose you add the same amount of heat to a 1.00 g sample of each of the following. Assuming each has the same initial temperature, which will have the highest final temperature? a. water (s = 4.184 J/(g·°C)) b. gasoline (s = 2.2 J/(g·°C)) c. sand (s = 0.799 J/(g·°C)) d. iron (s = 0.449 J/(g·°C)) e. gold (s = 0.129 J/(g·°C)) ANSWER: e 26. If 1,000 J of heat is applied to a 50.0 g iron bar, how much will its temperature increase? a. 10.0 ºC b. 22.2 ºC c. 44.4 ºC d. 50.0 ºC e. 100 ºC ANSWER: c 27. What mass of water (s = 4.184 J/(g·°C)) can absorb 5,000 J of heat while undergoing a temperature change of 25.0 °C? a. 18.0 g b. 23.9 g c. 41.9 g d. 47.8 g e. 83.8 g ANSWER: d 28. A 24.8 g sample of an unknown material was heated to 95.0 ºC and placed in 200 g of water (s = 4.184 J/g·°C) at 21.7 °C in a coffee-cup calorimeter. The final temperature of the material and water is 22.2 ºC. What is the identity of the material? a. aluminum, 0.904 J/g·°C b. zinc, 0.388 J/g·°C Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy c. tin, 0.227 J/g·°C d. silver, 0.235 J/g·°C e. lead, 0.127 J/g·°C ANSWER: d 29. A constant-pressure calorimeter can be used to determine _______, whereas a constant-volume calorimeter can be used to determine ___________. a. ΔH; ΔE b. ΔH; q c. w; ΔH d. ΔE; Δ H e. w; q ANSWER: a 30. What does the specific heat of a substance represent? a. The total energy of the substance at room temperature. b. The energy absorbed during a phase change. c. The heat lost by the substance in one second. d. The amount of heat required to raise 1 gram of the substance by 1 °C. e. The amount of heat required to change its temperature by 1 °C. ANSWER: d 31. In a calorimetry experiment, what does an increase in water temperature indicate? a. Heat is released from the surroundings. b. Heat is absorbed by the system. c. The system releases heat to the surroundings. d. The temperature difference is negligible. e. The calorimeter is malfunctioning.

ANSWER: c 32. If a system absorbs 500 J of heat but does no work, what happens to its internal energy? a. It increases by 500 J. b. It increases by 250 J. c. It decreases by 500 J. d. It decreases by 250 J. e. It remains unchanged.

ANSWER: a Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy 33. What is the main purpose of calorimetry? a. To measure temperature changes in the surroundings. b. To determine the mass of substances. c. To calculate energy changes during chemical or physical processes. d. To assess the thermal conductivity of materials. e. To compare specific heats of various substances.

ANSWER: c 34. A chemist places a 12.0-g sample of acetone (C3H6O) inside a bomb calorimeter with a heat capacity of 24.30 kJ/°C. When the acetone sample is combusted, the temperature of the calorimeter rises from 21.50 °C to 39.80 °C. How much heat did the acetone release per gram? a. 445 kJ/g b. 37.1 kJ/g c. –37.1 kJ/g d. 24.3 kJ/g e. 2.20 kJ/g ANSWER: b 35. What does the equation q = msΔT calculate? a. The amount of heat transferred during a temperature change. b. The heat released during a phase change. c. The energy stored in chemical bonds. d. The mass of a substance in a calorimeter. e. The total thermal conductivity of a material.

ANSWER: a 36. What assumption do you make about heat loss to the surroundings in calorimetry?

ANSWER We assume that the calorimeter is perfectly insulated, and that no heat is lost to : the surroundings.

37. A 100 g block of aluminum (s = 0.897 J/(g·°C)) is heated from 20.0 ºC to 70.0 ºC. How much heat is absorbed by the aluminum block?

ANSWER: 4490 J 38. A reaction in a calorimeter causes the temperature of 250 g of water (s = 4.184 J/(g·°C)) to increase by 12.5 ºC. How much heat is released by this reaction?

ANSWER: 13,100 J 39. Can two objects at the same temperature exchange heat? Why or why not?

ANSWER No, the temperature difference between two objects drives the transfer of heat, : so if there is no difference in temperature, there can be no exchange of heat.

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Chapter 6 Energy 40. A 37.5 g block of an unknown metal is heated to 95.0 ºC. It is then placed in a constant-pressure calorimeter which holds 100 g of water initially at 17.8 ºC. If the final temperature of the water and metal is 19.4 ºC, what is the specific heat capacity of the metal?

ANSWER: 0.236 J/g·ºC 41. Why do coastal areas have milder climates compared to inland areas?

ANSWER The high heat capacity of water can absorb and release large amounts of heat : with only small temperature changes, moderating the climate of coastal areas.

42. What is the difference between heat and enthalpy?

ANSWER Heat is the energy transferred between objects of different temperature. The : enthalpy combines the internal energy and PV work. Under constant pressure conditions, the two are equal.

43. In the combustion of toluene, what does ΔHrxn = −3910 kJ/mol mean? a. 3910 kJ is absorbed per mole of toluene burned. b. 3910 kJ is released per mole of toluene burned. c. The products have more energy than the reactants. d. The reaction is endothermic. e. The reaction involves no energy change.

ANSWER: b 44. How is ΔHrxn related to the coefficients in a balanced chemical equation? a. ΔHrxn is independent of the coefficients. b. ΔHrxn is only related to the products. c. ΔHrxn is only related to the reactants. d. ΔHrxn corresponds to the enthalpy change for the specified number of moles in the balanced equation. e. ΔHrxn is inversely proportional to the coefficients.

ANSWER: d 45. What happens to the surroundings in an exothermic reaction? a. The temperature remains constant. b. The surroundings absorb energy. c. The surroundings release energy. d. The temperature decreases. e. No energy is transferred.

ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy 46. Which phase change is associated with a negative enthalpy change? a. condensation b. melting c. sublimation d. vaporization e. none of the above ANSWER: a 47. Which of the following is NOT an extensive property? a. enthalpy b. heat c. internal energy d. volume e. temperature ANSWER: e 48. What is the enthalpy change when 2 moles of KOH dissolve in water, given that ΔHsoln = −57.7 kJ? a. –57.7 kJ b. –116 kJ c. –28.9 kJ d. –145 kJ e. 57.7 kJ ANSWER: b 49. What is the enthalpy change when 3.50 g of water (18.02 g/mol) is vaporized? ΔHvap = 44.0 kJ a. 8.55 kJ b. 154 kJ c. 0.0795 kJ d. 2780 kJ e. 1.43 kJ ANSWER: a 50. Potassium nitrate (KNO3), has an enthalpy of solution of 34.9 kJ. If 15.0 g of KNO3 is dissolved in 250.0 g of water with an initial temperature of 19.9 ºC, what is the final temperature of the water? a. 25.0 ºC b. 22.0 º Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy C c. 19.4 º C d. 17.0 º C e. 15.0 ºC ANSWER: e 51. What is the enthalpy change when 200 g of NaNO3 (85.00 g/mol) dissolves in water? ΔHsoln = 21.7 kJ a. 21.7 kJ b. 43.4 kJ c. 51.1 kJ d. 64.8 kJ e. 85.3 kJ ANSWER: c 52. How does the enthalpy of the products compare to the enthalpy of the reactants in an exothermic reaction?

ANSWER For an exothermic reaction, the enthalpy of the products is lower than the : enthalpy of the reactants. The difference is released as heat.

53. The neutralization of KOH with HNO3 is shown in the following reaction.

HNO3(aq) + KOH(aq) → KNO3(aq) + H2O(l) ΔHrxn = −56.2 kJ If 5.30 g of KOH is neutralized with excess nitric acid, how much heat is released?

ANSWER: –5.31 kJ

54. What is the heat change when 4.56 mol Cl2 reacts with excess Sn? Sn(s) + 2Cl2(g) → SnCl4(s) ΔHrxn = −511.3 kJ ANSWER: –1170 kJ 55. How much heat needs to be absorbed by 250 g of water to convert it from a liquid to a gas? ΔHvap = 44.0 kJ ANSWER: 611 kJ 56. Which of the following is a state function? a. work b. heat c. enthalp y d. time Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy e. distance ANSWER: c 57. What is the defining quality of a state function? a. It depends only on the path taken. b. It depends only on the initial and final states. c. It changes based on the surroundings d. It is always constant. e. It is always in the same physical state.

ANSWER: b 58. What does Hess’s law state? a. Enthalpy is inversely proportional to pressure. b. Enthalpy depends on the pathway. c. Heat and work are interchangeable. d. The total enthalpy change for a reaction is the sum of the enthalpy changes for the steps. e. Enthalpy is always positive.

ANSWER: d 59. What does a standard enthalpy of formation (ΔHf∘) of zero indicate? a. The substance is stable. b. The substance releases energy when formed. c. The substance is in its standard elemental form. d. The substance has no energy. e. The substance does not participate in reactions.

ANSWER: c 60. Which of the following describes how to determine the enthalpy of a reaction (ΔHrxn) from standard enthalpies of formation (ΔHf∘)? a. Subtract the product enthalpies of formation from the reactant enthalpies of formation. b. Subtract the reactant enthalpies of formation from the product enthalpies of formation. c. Add the product enthalpies of formation to the reactant enthalpies of formation. d. Multiply the product enthalpies of formation by the reactant enthalpies of formation. e. Divide the product enthalpies of formation by the reactant enthalpies of formation.

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Chapter 6 Energy ANSWER: b 61. How is a state function affected by the path of a reaction? a. It changes based on the path. b. It is constant along the path. c. It depends on heat transfer. d. It is independent of the path. e. It depends on work done.

ANSWER: d 62. What is the ∆Hrxn for the decomposition of hydrogen peroxide?

2H2O2(l) → 2H2O(l) + O2(g) ΔHf∘(H2O2(l)) = −187.8 kJ/mol ΔHf∘(H2O(l)) = −285.8 kJ/mol a. −196.0 kJ b. +196.0 kJ c. −759.4 kJ d. −375.6 kJ e. +375.6 kJ ANSWER: a 63. What is the ∆Hrxn for the formation of glucose from carbon dioxide and water?

6CO2(g) + 6H2O(l) → C6H12O6(s) + 6O2(g) ΔHf∘(CO2(g)) = −393.5 kJ/mol ΔHf∘(H2O(l)) = −285.8 kJ/mol ΔHf∘ (C6H12O6(s)) = −1273.0 kJ/mol a. −2802.8 kJ b. +2802.8 kJ c. −4075.8 kJ d. +1273.0 kJ e. −1273.0 kJ ANSWER: b 64. What is the ΔHrxn for the production of ammonia?

N2(g) + 3H2(g) → 2NH3(g) ΔHf∘(NH3(g)) = −45.9 kJ/mol a. −45.9 kJ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy b. +45.9 kJ c. −91.8 kJ d. +91.8 kJ e. 0 kJ ANSWER: c 65. If ΔH1 = −200 kJ and ΔH2 = +300 kJ, what is ΔH for the reaction that is the sum of the two reactions? a. −500 kJ b. −100 kJ c. 0 kJ d. +100 kJ e. +500 kJ ANSWER: d 66. What is required to calculate the enthalpy of a reaction using Hess's law? a. The temperature of each step of the reaction. b. The initial and final pressure. c. The amount of heat absorbed. d. The molar masses of the reactants. e. The enthalpy changes for each step of the reaction.

ANSWER: e 67. Which of the following processes does NOT involve a state function? a. heat transfer b. melting c. condensation d. sublimation e. vaporization ANSWER: a 68. What is the standard enthalpy of formation?

ANSWER It is the enthalpy change when one mole of a compound forms from its : elements in their standard states.

69. What is the difference between state functions, such as enthalpy and internal energy, and path functions, such as heat and work?

ANSWER State functions only depend on the state they are in. Path functions depend on : the path between two states.

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Chapter 6 Energy 70. Calculate the ΔHrxn for the neutralization of sodium hydroxide with hydrochloric acid.

NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) ΔHfo(NaOH(aq)) = −469.15 kJ/mol ΔHfo(HCl(aq)) = −166.106 kJ/mol ΔHfo(NaCl(aq)) = −407.27 kJ/mol ΔHfo(H2O(l)) = −285.83 kJ/mol ANSWER: –57.8 kJ 71. Calculate the ΔHrxn for the formation of carbon monoxide (CO) from its elements C(s) + O2(g) → CO(g) Given the following reactions and their enthalpy changes.

C(s) + O2(g) → CO2(g) ΔH = −393.5kJ CO(g) + O2(g) → CO2(g) ΔH = −283.0 kJ ANSWER: –110.5 kJ 72. Calculate ΔHrxn for the following reaction.

4NH3(g) + 5O2(g) → 6H2O(g) + 4NO(g) ΔHfo(NH3(g)) = −45.9 kJ/mol ΔHfo(NO(g)) = 90.3 kJ/mol ΔHfo(H2O(g)) = −241.8 kJ/mol ANSWER: –906.0 kJ 73. Calculate ΔHrxn for the combustion of benzene (C6H6).

C6H6(l) + O2(g) → 6CO2(g) + 3H2O(l)

ΔHfo(C6H6(l)) = −49.0 kJ/mol ΔHfo(CO2(g)) = −393.5 kJ/mol ΔHfo(H2O(l)) = −285.8 kJ/mol ANSWER: –3267.4 kJ 74. Which of the following reactions represents the standard enthalpy of formation of water? a. H2O(l) → H2(g) + O2(g) b. 2H2(g) + O2(g) → 2H2O(l) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 6 Energy c. H2(g) + O2(g) → H2O2(l) d. H2O(g) → H2O(l) e. 2H(g) + O(g) → H2O(l) ANSWER: b 75. What is the ΔHrxn for the decomposition of dinitrogen pentoxide?

2N2O5(g) → 4NO2(g) + O2(g) ΔHf∘(N2O5(g)) = +11.3 kJ/mol ΔHf∘ (NO2(g)) = +33.2 kJ/mol a. +88.2 kJ b. –22.6 kJ c. +110.2 kJ d. –21.9 kJ e. +44.5 kJ ANSWER: c

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Chapter 5 Stoichiometry 1. What is the mass of 2.50 moles of sodium chloride? a. 80.7 g b. 58.5 g c. 117 g d. 146 g e. 169 g ANSWER: d 2. How many moles are in 24 g of O2 gas? a. 0.50 moles b. 0.75 moles c. 1.0 mole d. 1.5 moles e. 2.0 moles ANSWER: b 3. How many molecules of nitrogen are in 0.120 moles of nitrogen gas? a. 7.23×1022 molecules b. 1.45×1023 molecules c. 6.02×1023 molecules d. 7.23×1023 molecules e. 3.61×1022 molecules ANSWER: a 4. A sample of water has a mass of 33.28 g. How many water molecules are in this sample? a. 1.112×1023 molecules b. 2.004×1024 molecules c. 1.112×1024 molecules d. 3.61×1026 molecules e. 1.112×1025 molecules ANSWER: d 5. A sample of acetone contains 4.59×1023 molecules of acetone. What is the mass of this acetone sample? The molar mass of acetone is 58.1 g/mol. a. 4.43 g b. 23.7 g Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry c. 44.3 g d. 58.1 g e. 0.0131 g ANSWER: c 6. What is the mass of 24.20 moles of carbon dioxide? The molar mass of CO2 is 44.01 g/mol. a. 978.0 g b. 1065 g c. 1068 g d. 1152 g e. 1357 g ANSWER: b 7. How many moles of Br are in 3.84 moles of MgBr2? a. 2.00 moles b. 3.00 moles c. 3.84 moles d. 7.68 moles e. 11.5 moles ANSWER: d 8. A sample contains 2.5×1024 atoms of silver (Ag). How many moles of silver are in the sample? a. 0.42 moles b. 1.2 moles c. 1.8 moles d. 2.4 moles e. 4.2 moles ANSWER: d 9. How many sulfur atoms are in a 53.2 g sample of CS2? a. 7.88×1023 atoms b. 8.42×1024 atoms c. 1.53×1024 atoms d. 4.21×1023 atoms Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry e. 8.42×1023 atoms ANSWER: e 10. What is the mass of 9.03×1023 Ar atoms? a. 59.9 g b. 39.7 g c. 17.3 g d. 6.00 g e. 120 g ANSWER: a 11. How many molecules are in a 54.9 g sample of carbon monoxide (CO)?

ANSWER: 1.18×1024 molecules 12. How many moles of CuSO4 are in a 135.61 g sample?

ANSWER: 0.84963 mol 13. Hydrochloric acid (HCl) has a molar mass of 36.46 g/mol. How many molecules are in 81.00 g of HCl? ANSWER: 1.338×1024 molecules 14. How many fluorine atoms are in 44.31 g of BF3?

ANSWER: 1.181×1024 F atoms 15. What is the mass of 1.99×1025 GeCl2 molecules?

ANSWER: 4740 g 16. How many moles of HNO3 can be produced from 5.0 moles of NO2? a. 2.5 moles b. 3.3 moles c. 5.0 moles d. 6.7 moles e. 10 moles ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry 17. Calcium carbonate can decompose according to the following reaction. How many grams of CaO are formed from the decomposition of 35.2 g of CaCO3? a. 15.8 g b. 21.5 g c. 25.3 g d. 19.7 g e. 30.6 g ANSWER: d 18. If 62.4 g of BaSO4 is produced from the following reaction, what mass of NaCl is also formed? a. 15.7 g b. 31.2 g c. 45.6 g d. 48.9 g e. 62.6 g ANSWER: b 19. Ammonium nitrate decomposes via the following reaction. If 10 g of ammonium nitrate decomposes, what mass of water is produced? a. 2.0 g b. 3.6 g c. 4.5 g d. 5.8 g e. 7.2 g ANSWER: c 20. The combustion of butane is shown in the following reaction. If 3.65 g of butane is combusted, what mass of carbon dioxide is produced? a. 8.52 g b. 9.76 g c. 22.1 g d. 11.1 g e. 5.70 g ANSWER: d 21. What is the first step when calculating the mass of a product produced from a given mass of a reactant?

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Chapter 5 Stoichiometry a. Use the molar mass of the product to convert from moles to grams. b. Use the coefficients in the balanced equation to convert from moles of reactant to moles of product. c. Use the molar mass of the reactant to convert from grams to moles. d. Use the coefficients in the balanced equation to convert from moles of product to moles of reactant. e. Use the molar mass of the product to convert from grams to moles.

ANSWER: c 22. Sulfur hexafluoride is formed in the following reaction. How many moles of SF6 can be formed from 41.9 g of F2?

S + 3F2 → SF6 a. 1.10 mol SF6 b. 0.368 mol SF6 c. 0.550 mol SF6 d. 0.730 mol SF6 e. 3.31 mol SF6 ANSWER: b 23. Calcium hydroxide and phosphoric acid react according to the following acid-base reaction. What mass of water is produced from 15.2 g of H3PO4?

3Ca(OH)2 + 2H3PO4 → Ca3(PO4)2 + 6H2O a. 5.51 g b. 16.8 g c. 8.38 g d. 2.80 g e. 9.33 g ANSWER: c 24. What mass of copper is produced when 33.7 g of iron reacts with excess Cu2SO4?

Fe + Cu2SO4 → 2Cu + FeSO4 a. 50.9 g b. 63.5 g c. 88.5 g d. 76.7 g e. 103 g ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry 25. What mass of Pb(NO3)2 should be used to produce 100 g of KNO3 through the following reaction?

Pb(NO3)2 + 2KI → PbI2 + 2KNO3 a. 90.3 g b. 328 g c. 100 g d. 81.9 g e. 164 g ANSWER: e 26. If 3 moles of C6H12 are combusted in the following reaction, how many moles of water are produced?

C6H12 + 9O2 → 6CO2 + 6H2O ANSWER: 18 moles 27. Elemental magnesium reacts with hydrochloric acid according to the following reaction.

Mg + 2HCl → MgCl2 + H2 If a 5.00 g piece of Mg is added to a beaker containing HCl, what mass of MgCl2 is produced?

ANSWER: 19.6 g 28. When heated, potassium chlorate decomposes according to the following reaction.

2KClO3 → 2KCl + 3O2 How many grams of KClO3 were heated if 0.11 moles of O2 are produced?

ANSWER: 9.0 g 29. What mass of K3PO4 is produced when 47.2 g of KOH is reacted with excess H3PO4?

H3PO4 + 3KOH → K3PO4 + 3H2O ANSWER: 59.5 g 30. Phosphorus and chlorine react according to the following reaction.

P4 + 10Cl2 → 4PCl5 What mass of phosphorus is needed to produce 45.4 g of PCl5?

ANSWER: 6.75 g 31. Cesium and bromine react to form cesium bromide according to the following equation. If 3 moles of Cs combine with 2 moles of Br2, how many moles of CsBr are formed?

2Cs + Br2 → 2CsBr a. 3 moles b. 2 moles c. 1.5 moles Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry d. 7 moles e. 4 moles ANSWER: a 32. Sodium and chlorine react to form sodium chloride according to the following equation. If 3 moles of Na combine with 2 moles of Cl2, how many moles of NaCl are formed?

2Na + Cl2 → 2NaCl a. 3 moles b. 2 moles c. 1 mole d. 6 moles e. 4 moles ANSWER: e 33. Propane (C3H8) undergoes the following combustion reaction. If 15 g of propane is combusted with 15 g of oxygen, what mass of water is produced?

C3H8 + 5O2 → 3CO2 + 4H2O a. 24.5 g b. 6.12 g c. 3.38 g d. 10.6 g e. 6.76 g ANSWER: e 34. Cl2 reacts with NaBr according to the following reaction. If 2.04 g of NaBr react with 1.43 g of Cl2, what mass of Br2 is produced?

Cl2 + 2NaBr → 2NaCl + Br2 a. 0.790 g b. 1.02 g c. 1.58 g d. 2.04 g e. 3.22 g ANSWER: c 35. What is the limiting reagent when 1.5 moles of carbonic acid are added to 2.0 moles of NaOH? How many moles of sodium carbonate are formed in this reaction?

H2CO3 + 2NaOH → Na2CO3 + 2H2O a. NaOH; 1.0 mol Na2CO3 b. NaOH; 1.5 mol Na2CO3 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry c. NaOH; 3.0 mol Na2CO3 d. H2CO3; 1.0 mol Na2CO3 e. H2CO3; 1.5 mol Na2CO3 ANSWER: a 36. When heated, iron(III)oxide reacts with carbon monoxide according to the following equation. If 37.0 g of Fe2O3 reacts with 18.0 g of CO, what mass of iron is produced?

Fe2O3 + 3CO → 2Fe + 3CO2 a. 25.9 g b. 18.5 g c. 20.5 g d. 23.9 g e. 26.3 g ANSWER: d 37. Acetylene reacts with oxygen according to the following equation. If 1.5 moles of acetylene react with 3.0 moles of oxygen, what reactant is left over and in what amount?

2C2H2 + 5O2 → 4CO2 + 2H2O a. acetylene; 0.30 moles b. acetylene; 0.50 moles c. acetylene; 3.0 moles d. oxygen; 3.0 moles e. oxygen; 1.5 moles ANSWER: a 38. Titanium(IV) oxide is produced by the following reaction. If 100 g of FeTiO3 reacts with 55.0 g of H2SO4, what reactant is left over and in what amount?

FeTiO3 + H2SO4 → TiO2 + FeSO4 + H2O a. 29.7 g FeTiO3 b. 14.9 g FeTiO3 c. 6.54 g H2SO4 d. 13.3 g H2SO4 e. Neither reactant is in excess.

ANSWER: b 39. Carbon dioxide reacts with barium oxide to form barium carbonate, BaCO3. If 0.230 moles of CO2 react with 0.870 moles of BaO, what is the limiting reagent and how much barium carbonate is produced? a. CO2; 22.7 g Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry b. CO2; 45.4 g c. BaO; 45.4 g d. BaO; 90.8 g e. BaO; 22.7 g ANSWER: b 40. In the process of purifying nickel, the following reaction occurs. If 0.790 moles of Ni reacts with 2.31 moles of CO, what mass of Ni(CO)4 is produced?

Ni + 4CO → Ni(CO)4 a. 135 g b. 39.4 g c. 58.7 g d. 98.6 g e. 171 g ANSWER: d 41. If 3.44 moles of H2S reacts with 1.54 moles of SO2, what is the limiting reagent and how many moles of sulfur are produced?

2H2S + SO2 → 3S + 2H2O ANSWER: SO2 is the limiting reagent and 4.62 moles of S are produced.

42. Aluminum reacts with sulfuric acid via the following reaction.

2Al + 3H2SO4 → Al2(SO4)3 + 3H2 If 3.20 g of aluminum reacts with 17.1 g of sulfuric acid, what is the limiting reagent and what mass of aluminum sulfate is produced?

ANSWER: Sulfuric acid is the limiting reagent and 19.9 g of aluminum sulfate are produced.

43. Barium chloride reacts with sodium sulfate via the following reaction.

BaCl2 + Na2SO4 → 2NaCl + BaSO4 If 32.5 g of BaCl2 reacts with 32.0 g of Na2SO4, which reactant is in excess and how much of the reactant is left after the reaction goes to completion?

ANSWER Na2SO4 is in excess and there is 9.8 g of it remaining after the reaction goes to : completion.

44. If 24.97 g of Na2S2O3 reacts with 17.90 g of I2, what is the limiting reagent and what mass of NaI is produced?

2Na2S2O3 + I2 → Na2S4O6 + 2NaI ANSWER: Iodine is the limiting reagent and 21.14 g of NaI are produced.

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Chapter 5 Stoichiometry 45. If 15.5 g of FeCl3 react with 8.2 g of H2S, what is the limiting reagent and what mass of sulfur is produced?

2FeCl3 + H2S → 2FeCl2 + 2HCl + S ANSWER: FeCl3 is the limiting reagent and 1.5 g of S are produced.

46. The major component of a solution is called the a. solvent. b. solute. c. solution. d. aqueous fraction. e. major fraction.

ANSWER: a 47. What is the definition of molarity? a. moles of solute per kilogram of solvent b. moles of solute per liter of solution c. moles of solute per kilogram of solution d. moles of solute per liter of solvent e. grams of solute per milliliter of solution ANSWER: b 48. How many moles of CuSO4 are in 500 mL of a 2.1 M copper(II) sulfate solution? a. 4.2 moles b. 2.1 moles c. 1.1 moles d. 0.52 moles e. 0.25 moles ANSWER: c 49. How many grams of glucose (C6H12O6) are needed to prepare 250 mL of a 0.500 M glucose solution? a. 90.1 g b. 180 g c. 45.0 g d. 225 g e. 22.5 g ANSWER: e 50. What is the molarity of a solution prepared by dissolving 37.2 g of KOH in enough water to make 0.500 L of solution?

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Chapter 5 Stoichiometry a. 0.133 M b. 0.660 M c. 1.33 M d. 3.32 M e. None of the above.

ANSWER: c 51. What is the last step when making a solution using a volumetric flask? a. Add the solid solute. b. Partially fill the flask with solvent. c. Add solvent in an amount equal to the final volume of the solution. d. Dilute the solution to the mark on the flask with solvent. e. Dilute the solution to the top of the flask with solvent.

ANSWER: d 52. When diluting a solution with additional solvent, the amount of moles of solute in the dilute solution a. is greater than the amount of moles of solute in the original solution. b. is less than the amount of moles of solute in the original solution. c. is the same as the amount of moles of solute in the original solution. d. relative to the original solution cannot be determined without more information. e. relative to the original solution depends on the solute.

ANSWER: c 53. What volume of a 3.00 M KNO3 solution would you need to make 250 mL of a 1.43 M solution? a. 15.1 mL b. 83.3 mL c. 65.6 mL d. 42.1 mL e. 119 mL ANSWER: e 54. Which of the following should be used for the gravimetric analysis of OH¯? a. Mg2+ b. Ba2+ c. Li+ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry d. Zn2+ e. K+ ANSWER: d 55. How many grams of Cu2SO4 are needed to make 100 mL of a 0.300 M solution of copper(I) sulfate? a. 6.69 g b. 22.3 g c. 13.4 g d. 0.670 g e. 4.46 g ANSWER: a 56. How many grams of ammonium cyanide (NH4CN) are needed to make 1.5 L of a 3.0 M solution?

ANSWER: 2.0×102 g 57. How many milliliters of a 3.40 M K2SO3 solution are needed to make 100 mL of a 0.500 M solution?

ANSWER: 14.7 mL 58. A 1.5 L solution contains 51.6 g of Mg(C2H3O2)2. What is the molarity of the solution?

ANSWER: 0.24 M 59. Cobalt(II) chloride hexahydrate (CoCl2·6H2O) turns from pink to blue when dehydrated. If you want to make 250 mL of a 0.33 M cobalt(II) chloride hexahydrate solution, how many grams would you need?

ANSWER: 19.6 g 60. 300.0 mL of a 0.600 M aluminum chloride (AlCl3) solution are mixed with 400.0 mL of a 0.800 M sodium hydroxide (NaOH) solution. Which reactant is the limiting reagent and how many grams of aluminum hydroxide are formed?

AlCl3(aq) + 3NaOH(aq) → Al(OH)3(s) + 3NaCl(aq) ANSWER: NaOH is the limiting reagent and 8.32 g of Al(OH)3 are produced.

61. The actual yield of a reaction is a. the experimentally measured amount of product formed. b. the maximum amount of product predicted. c. the amount of reactants used up. d. always equal to the theoretical yield. e. the reactant left unreacted.

ANSWER: a Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry 62. The theoretical yield of a reaction is a. the experimentally measured amount of product formed. b. the maximum amount of product predicted. c. the amount of reactants used up. d. always equal to the actual yield. e. the reactant left unreacted.

ANSWER: b 63. A reaction has a theoretical yield of 22.5 g and a percent yield of 73.1%. What is the actual yield? a. 1.64 g b. 16.4 g c. 8.20 g d. 22.5 g e. 15.8 g ANSWER: b 64. If the theoretical yield of a product is 15 g, but only 12 g is obtained experimentally, what is the percent yield? a. 67% b. 75% c. 80% d. 90% e. 100% ANSWER: c 65. If a chemist calculates a percent yield of 120%, what is the most likely explanation? a. The reaction was incomplete. b. The theoretical yield was overestimated. c. The reactants were impure. d. The reaction was carried out in perfect conditions. e. The actual yield was measured incorrectly.

ANSWER: e 66. Calcium nitrate reacts with sodium phosphate according to the following reaction.

3Ca(NO3)2(aq) + 2Na3PO4(aq) → Ca3(PO4)2(s) + 6NaNO3(aq) If 50.0 g of Ca(NO3)2 reacts with 40.0 g of Na3PO4, what is the theoretical yield of Ca3(PO4)2? a. 40.0 g b. 63.0 g c. 31.5 g Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry d. 25.9 g e. 3.15 g ANSWER: c 67. Sodium bicarbonate decomposes upon heating to form sodium carbonate, carbon dioxide, and water.

2NaHCO3(s) → Na2CO3(s) + CO2(g) + H2O(g) If 12.0 g of NaHCO3 is heated, what is the theoretical yield of sodium carbonate? a. 3.80 g b. 7.57 g c. 15.1 g d. 17.9 g e. 22.8 g ANSWER: b 68. Silver nitrate reacts with potassium bromide to form silver bromide.

AgNO3(aq) + KBr(aq) → AgBr(s) + KNO3(aq) If 8.71 g of AgNO3 reacts with excess KBr, what is the percent yield if 38.0 g of AgBr are produced? a. 8.30% b. 12.7% c. 21.6% d. 25.3% e. 28.0% ANSWER: d 69. Barium chloride reacts with sodium carbonate to form barium carbonate.

BaCl2(aq) + Na2CO3(aq) → BaCO3(s) + 2NaCl(aq) If 100 mL of 0.30 M BaCl2 react with excess sodium carbonate, what is the theoretical yield of BaCO3? a. 2.4 g b. 2.9 g c. 5.9 g d. 6.5 g e. 12 g ANSWER: c 70. Xenon tetrafluoride is formed by reacting xenon with fluorine.

Xe + 2F2 → XeF4 If the theoretical yield of XeF4 is 55.6 g, what is the percent yield if the actual yield is 22.1 g? a. 4.00% b. 13.2% Copyright Macmillan Learning. Powered by Cognero.

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Chapter 5 Stoichiometry c. 19.8% d. 39.7% e. 43.8% ANSWER: d 71. A student conducts a reaction and the actual amount of product obtained is 3.23 g. If the theoretical yield is 8.99 g, what is the percent yield for this reaction?

ANSWER: 35.9% 72. A factory that produces food flavorings in bulk makes a particular product with a percent yield of 86.30%. If the theoretical yield of a typical batch is 4.55 kg, what is the actual yield of a typical batch?

ANSWER: 3.93 kg 73. A common substitution reaction has a percent yield of 34.3%. If the actual yield from the reaction is 4.20 g, what is the theoretical yield of the reaction?

ANSWER: 12.2 g 74. Potassium chlorate decomposes according to the following reaction.

2KClO3(s) → 2KCl(s) + 3O2(g) What is the percent yield if 15 g of O2 are actually produced, but the reaction can theoretically produce 19.7 g of O2?

ANSWER: 76% 75. Magnesium reacts with hydrochloric acid according to the following reaction.

Mg(s) + 2HCl(aq) → MgCl2(s) + H2(g) If 2.00 g of Mg reacts with a solution containing 0.300 moles of HCl, what is the theoretical yield of MgCl2 in grams for this reaction?

ANSWER: 7.84 g

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Chapter 4 Chemical Reactions 1. Which set of coefficients would balance the following chemical equation? a. 1, 3, 2, 4 b. 2, 5, 6, 2 c. 1, 5, 3, 4 d. 2, 10, 6, 8 e. 1, 7, 3, 4 ANSWER: c 2. Which of the following is TRUE about the products of a combustion of a hydrocarbon? a. The products are always water and hydrogen. b. The products are always water and carbon monoxide. c. The products are always water and carbon dioxide. d. The products are always hydrogen and carbon dioxide. e. The products are always nitrogen and carbon dioxide.

ANSWER: c 3. What must be the same on both sides of a balanced chemical equation? a. the number of molecules b. the number of atoms of each element c. the volume of gases d. the total energy e. the temperature ANSWER: b 4. Which of the following represents a correctly balanced equation? a. b. c. d. e.

ANSWER: d 5. In the equation

, what is the ratio of chlorine molecules to aluminum chloride

molecules? a. 3:2 b. 1:1 c. 2:3 d. 1:2 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 4 Chemical Reactions e. 3:3 ANSWER: a 6. What is the coefficient for water when the following equation is balanced? a. 1 b. 2 c. 3 d. 4 e. 6 ANSWER: e 7. What is the coefficient for sodium hydroxide when the following equation is balanced? a. 1 b. 2 c. 3 d. 4 e. 6 ANSWER: c 8. What is the best way to balance a combustion reaction? a. adjusting only the hydrogen atoms b. using water as a reactant c. ensuring that the oxygen atoms are balanced first d. balancing carbon atoms first e. making sure the number of oxygen molecules is always even ANSWER: d 9. Balance the chemical equation for the combustion of nonane.

ANSWER: 10. Why do we not change the subscripts when balancing chemical equations?

ANSWER The subscripts tell us how many atoms are present in each molecule. If they are : changed, the identity of the molecule has changed, and the chemical reaction is different.

11. Balance the following chemical equation.

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Chapter 4 Chemical Reactions 12. Balance the reaction of dichlorine pentoxide with water to form chloric acid.

ANSWER: 13. Which of the following reaction types is not an oxidation-reduction reaction? a. metal-acid b. metal displacement c. combustion d. acid-base neutralization e. metal-nonmetal ANSWER: d 14. Which of the following is a synthesis reaction? a. b. c. d. e.

ANSWER: a 15. Which type of reaction is represented by

?

a. synthesis b. decomposition c. single-displacement d. doubledisplacement e. combustion ANSWER: b 16. Which of the following is a single-displacement reaction? a. b. c. d. e.

ANSWER: c 17. Which element is oxidized in the following reaction?

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Chapter 4 Chemical Reactions a. aluminum b. iron c. oxygen d. both aluminum and iron e. This is not an oxidation-reduction reaction.

ANSWER: a 18. In a single-displacement reaction involving a metal and a nonmetal, which species is typically reduced? a. the metal b. the nonmetal c. both are reduced d. neither is reduced e. the product is reduced ANSWER: b 19. When a metal reacts with an acid (H+), what will the products be? a. a cation and a metal b. an anion and a metal c. a cation and hydrogen gas d. an anion and hydrogen gas e. a metal and hydrogen gas ANSWER: c 20. In a single-displacement reaction involving a metal and a nonmetal, which species is typically oxidized? a. the metal b. the nonmetal c. both are reduced d. neither is reduced e. the product is reduced ANSWER: a 21. Write a balanced equation that describes the combustion of calcium.

ANSWER: 22. What is the difference between a synthesis reaction and a decomposition reaction?

ANSWER A synthesis reaction combines multiple elements into a single compound, : while a decomposition reaction breaks down a compound into its constituent elements.

23. Describe the difference between a single-displacement reaction and a double-displacement reaction.

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Chapter 4 Chemical Reactions ANSWER A single-displacement reaction is usually between a metal and an ionic : compound, and is also a redox reaction, whereas a double-displacement reaction is the swapping of ions between two ionic compounds.

24. Write a balanced equation for the reaction of zinc with nitrogen.

ANSWER: 25. What is the ionic equation when sodium sulfate is dissolved in water? a. Na(SO4)3(s) → 2Na+(aq) + 3SO42–(aq) b. Na3SO4(s) → 3Na+(aq) + SO42–(aq) c. Na(SO4)2(s) → 2Na+(aq) + 2SO42–(aq) d. Na2SO4(s) → 2Na+(aq) + SO42–(aq) e. NaSO4(s) → Na+(aq) + SO42–(aq) ANSWER: d 26. What happens when ammonium chloride dissolves in water? a. It remains a solid. b. It dissociates into NH4+ and Cl– ions. c. It reacts with water to form HCl. d. It decomposes into ammonia and chlorine gas. e. It forms a solid.

ANSWER: b 27. What is the ionic equation when sodium sulfide is dissolved in water? a. b. c. d. e.

ANSWER: c 28. What is the ionic equation when aluminum sulfate is dissolved in water? a. Al2(SO4)3(s) → 2Al3+(aq) + 3SO42–(aq) b. Al3(SO4)2(s) → 3Al2+(aq) + 2SO43–(aq) c. AlSO4(s) → Al2+(aq) + SO42–(aq) d. AlSO4(s) → Al3+(aq) + SO43–(aq) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 4 Chemical Reactions e. Al(SO4)3(s) → Al3+(aq) + 3SO4–(aq) ANSWER: a 29. When potassium bicarbonate is dissolved in water, how many ions are produced? a. 0 b. 1 c. 2 d. 3 e. 4 ANSWER: c 30. What is the ionic equation when barium iodide is dissolved in water? a. b. c. d. e.

ANSWER: b 31. What is the balanced ionic equation when solid ammonium bromide is dissolved in water? a. b. c. d. e. none of the above ANSWER: d 32. What are electrolyte solutions? a. All solutions where water is the solvent. b. Solutions that have a net negative charge. c. Solutions that have a net positive charge. d. Solutions connected to an electrical circuit. e. Solutions containing dissolved ions.

ANSWER: e 33. What is the ionic equation when solid calcium chloride is dissolved in water? a. CaCl(s) → Ca+(aq) + Cl–(aq) Copyright Macmillan Learning. Powered by Cognero.

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Chapter 4 Chemical Reactions b. Ca2Cl(s) → 2Ca+(aq) + Cl2–(aq) c. CaCl(s) → Ca2+(aq) + Cl2–(aq) d. CaCl2(s) → Ca2+(aq) + 2Cl–(aq) e. CaCl3(s) → Ca3+(aq) + 3Cl–(aq) ANSWER: d 34. Write the ionic equation for the dissolution of KC2H3O2 in water.

ANSWER: KC2H3O2(s) → K+(aq) + C2H3O2–(aq) 35. Write the ionic equation for the dissolution of aluminum bromide in water.

ANSWER: AlBr3(s) → Al3+(aq) + 3Br–(aq) 36. Write the molecular and ionic equations for the dissolution of calcium nitrate in water.

ANSWER: molecular: Ca(NO3)2(s) → Ca(NO3)2(aq) ionic: Ca(NO3)2(s) → Ca2+(aq) + 2NO3–(aq) 37. Which of the following compounds is soluble in water? a. AgCl b. PbI2 c. BaSO4 d. KNO3 e. MgCO3 ANSWER: d 38. Which of the following compounds is insoluble in water? a. Hg(C2H3O2)2 b. AgClO4 c. Li2SO3 d. Ba(OH)2 e. Cu3(PO4)2 ANSWER: e 39. A complete ionic equation shows a. positive ions. b. negative ions. c. spectator ions. d. precipitated species.

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Chapter 4 Chemical Reactions e. all of the above.

ANSWER: e 40. In which of the following aqueous combinations would no precipitation reaction occur? a. LiCl + K3PO4 b. NaSO4 + KNO3 c. NH4OH + PbNO3 d. AgClO4 + CaBr2 e. Fe2S3 + Hg(C2H3O2)2 ANSWER: b 41. In which of the following aqueous combinations would a precipitation reaction occur? a. LiCl + (NH4)2CO3 b. Na2SO4 + KNO3 c. NaC2H3O2 + CuSO4 d. AgClO4 + CaBr2 e. ZnS + AlCl3 ANSWER: d 42. Which of the following compounds would cause a precipitate to form when added to a solution of Ba2+? a. NH4OH b. NaC2H3O2 c. MgBr2 d. Na2SO4 e. FeI3 ANSWER: d 43. Which of the following will precipitate in an aqueous solution? a. Mg(C2H3O2)2 b. NH4ClO4 c. LiI d. KOH e. Ba (PO ) 3

4 2

ANSWER: e 44. When the temperature of an aqueous solution is increased from room temperature to just below boiling, the Copyright Macmillan Learning. Powered by Cognero.

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Chapter 4 Chemical Reactions solubility of ionic compounds in the solution generally a. increases greatly. b. decreases greatly. c. stays about the same. d. is always constant for all solutes in aqueous solutions. e. cannot determine without more information.

ANSWER: a 45. What species will precipitate when BaCl2 and Na2SO4 are mixed?

ANSWER: BaSO4 46. What species will precipitate when ZnS(aq) and K3PO4(aq) are mixed?

ANSWER: Zn3(PO4)2 47. Will a precipitate form when CaCl2 is mixed with Fe2S3, and if so, what is the precipitate?

ANSWER: No, there will not be a precipitate formed. CaS and FeCl3 are both soluble.

48. Will a precipitate be formed when Hg(NO3)2 is mixed with MgI2, and if so, what is the precipitate?

ANSWER: Yes, the precipitate is HgI2.

49. Identify the precipitate formed when Mg(NO3)2 reacts with K3PO4, and write the balanced net ionic equation for the reaction.

ANSWER: The precipitate is Mg3(PO4)2.

3Mg2+(aq) + 2PO43– (aq)

Mg3(PO4)2(s)

50. When nitric acid reacts with potassium hydroxide, the products will be water and what? a. KOH b. NaNO3 c. KNO3 d. NaCl e. HNO3 ANSWER: c 51. What ions will be present in an aqueous solution of sodium hydroxide? a. Na+ and Cl– b. OH- and Cl– c. H+ and OH– Copyright Macmillan Learning. Powered by Cognero.

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Chapter 4 Chemical Reactions d. Na+ and OH– e. Na+ + H2 ANSWER: d 52. Which of the following is the net ionic equation for the neutralization of HCl and NaOH? a. b. c. d. e.

ANSWER: b 53. When an acid reacts with a base, they form water and what other kind of compound? a. a salt b. a gas c. an acid d. a base e. It depends on what the acid and base are.

ANSWER: a 54. Which of the following is the net ionic equation for the neutralization of HNO3 and KOH? a. b. c. d. e.

ANSWER: b 55. Predict the products for the following reaction.

H3PO4(aq) + Ba(OH)2(aq) → a. Ba3PO4(s) + H2O(l) b. BaPO4(s) + H2O(l) c. Ba3(PO4)2(s) + H2O(l) d. Ba3(PO4)2(s) + H2O(aq) e. none of the above Copyright Macmillan Learning. Powered by Cognero.

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Chapter 4 Chemical Reactions ANSWER: c 56. When HI reacts with Ba(OH)2, how many molecules of water are produced? a. 0 H2O b. 1 H2O c. 2 H2O d. 4 H2O e. 8 H2O ANSWER: c 57. How many sodium hydroxide molecules are needed to form sodium phosphate when phosphoric acid reacts with sodium hydroxide? a. 0 NaOH b. 1 NaO H c. 2 NaO H d. 3 NaOH e. 6 NaOH ANSWER: d 58. What is the driving force when an acid-base neutralization reaction occurs?

ANSWER: the formation of water 59. What is the net ionic equation of an acid-base neutralization reaction?

ANSWER: H+(aq) + OH–(aq) → H2O(l) 60. Identify the products, and write the balanced molecular equation for the neutralization reaction between H2SO4 and KOH. ANSWER The products are K2SO4(aq) and H2O(l). The balanced reaction is H2SO4(aq) + : 2KOH(aq) → 2H2O(l) + K2SO4(aq).

61. Identify the products, and write the balanced molecular equation for the acid-base neutralization reaction between HClO4 and LiOH. ANSWER The products are LiClO4(aq) and H2O(l). The balanced reaction is HClO4(aq) : + LiOH(aq) → H2O(l) + LiClO4(aq).

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Chapter 4 Chemical Reactions 62. What is a base?

ANSWER: A base is a substance that produces OH– in aqueous solution.

63. According to the activity series, which of the following metals can displace hydrogen from an acid? a. copper b. zinc c. silver d. gold e. platinu m ANSWER: b 64. Which of the following metals will react most strongly with water? a. calcium b. magnesiu m c. zinc d. potassium e. copper ANSWER: d 65. What happens when magnesium is placed in hydrochloric acid? a. No reaction occurs. b. Magnesium hydride and chlorine gas are formed. c. Water and magnesium hydroxide are formed. d. Hydrogen gas and magnesium chloride are formed. e. Magnesium oxide is formed.

ANSWER: d 66. Which of the following metals will NOT react with water at room temperature? a. lithium b. sodium c. copper d. calcium e. potassium ANSWER: c 67. Which of the following metals is hardest to oxidize? a. silver b. iron Copyright Macmillan Learning. Powered by Cognero.

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Chapter 4 Chemical Reactions c. copper d. nickel e. gold ANSWER: e 68. Which group of metals is generally the most reactive? a. alkali metals b. alkaline earth metals c. transition metals d. metalloids e. All metals have about the same reactivity.

ANSWER: a 69. Which of the following is TRUE about the activity series? a. An element can reduce an ion above it on the activity series. b. An ion can oxidize an element above it on the activity series. c. An ion can reduce an element below it on the activity series. d. An element can reduce an ion below it on the activity series. e. Cannot tell without more information.

ANSWER: d 70. What is the balanced equation when barium reacts with steam? a. b. c. d. e.

ANSWER: b 71. Based on the activity series, which of the following metals can reduce cobalt(II) to elemental cobalt? a. nickel b. tin c. copper d. lead e. none of the above ANSWER: e 72. Write the net ionic equation for the reaction of tin with silver nitrate.

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Chapter 4 Chemical Reactions ANSWER: Sn(s) + 2Ag+(aq) → Sn2+(aq) + 2Ag(s) 73. Complete and balance the following reaction. If a reaction doesn't occur, write no reaction.

ANSWER: no reaction 74. Write the molecular, total, and net ionic equations for the reaction of Pb with CuSO4.

ANSWER: molecular: Pb(s) + CuSO4(aq) → PbSO4(aq) + Cu(s) ionic: Pb(s) + Cu2+(aq) + SO42–(aq) → Pb2+(aq) + SO42–(aq) + Cu(s) net ionic: Pb(s) + Cu2+(aq) → Pb2+(aq) + Cu(s) 75. Complete and balance the following reaction.

ANSWER: Ca + H2SO4 → CaSO4 + H2

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Chapter 3 From Elements to Compounds 1. What is the net charge of an ion with 7 protons, 8 neutrons, and 4 electrons? a. –4 b. –1 c. + 1 d. + 2 e. + 3 ANSWER: e 2. What charge would you expect on an ion of strontium? a. + 4 b. + 2 c. + 1 d. –1 e. –2 ANSWER: b 3. Which of the following anions has the most oxygen atoms? a. perchlorate b. bromate c. chlorite d. hypobromite e. iodide ANSWER: a 4. What is the formula for the nitrate ion? a. NO32– b. NO3– c. NO4– d. NO42– e. N3– ANSWER: b 5. Which of the following metals commonly forms an ion with a +3 charge? a. zinc Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds b. copper c. magnesiu m d. sodium e. aluminum ANSWER: e 6. What is the net charge of an ion with 12 protons, 11 neutrons, and 10 electrons? a. + 2 b. + 1 c. –1 d. –2 e. –3 ANSWER: a 7. Based on its position in the periodic table, what charge would you expect an ion of phosphorus to have? a. + 2 b. + 3 c. –3 d. –2 e. –1 ANSWER: c 8. Which of the following is the ferrous ion? a. Fe3– b. Fe+ c. Fe2+ d. Fe3+ e. Fe2– ANSWER: c 9. Which of the following does NOT commonly form a monatomic anion? a. xenon b. phosphorus c. bromine Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds d. sulfur e. nitrogen ANSWER: a 10. If an ion has 2 more protons than neutrons, what will the charge be on that ion? a. –2 b. +1 c. +2 d. +4 e. We cannot know with only the given information.

ANSWER: e 11. What is the charge on the carbonate ion? a. + 1 b. –1 c. + 2 d. –2 e. –3 ANSWER: d 12. Which polyatomic anion has the formula HCO3–?

ANSWER: The bicarbonate anion or the hydrogen carbonate anion.

13. What is the formula for the ammonium ion?

ANSWER: NH4+ 14. If a neutral atom loses an electron, what will the charge be on the newly formed ion?

ANSWER: +1 15. Which common polyatomic ion has a charge of –3?

ANSWER: phosphate, PO43– 16. Based on its position in the periodic table, what would you expect the charge on an ion of iodine to be?

ANSWER: –1 17. What is the formula of an ionic compound made from iron(III) and fluoride ions? a. FeF b. Fe3F Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds c. Fe2F3 d. FeF3 e. Fe3F2 ANSWER: d 18. What is the formula for barium sulfate? a. BaSO3 b. BaSO4 c. Ba2SO4 d. Ba(SO4)2 e. Ba(SO3)2 ANSWER: b 19. What is the name of the ionic compound with the formula FeO? a. iron(II) oxide b. iron oxide c. iron(III) oxide d. iron dioxide e. iron peroxide ANSWER: a 20. What is the formula of calcium phosphate? a. Ca3PO4 b. Ca3(PO4)2 c. Ca2(PO4)3 d. CaPO4 e. Ca2PO4 ANSWER: b 21. What is the formula of cobalt(II) chloride hexahydrate? a. CuCl2 4H2O b. CoCl2 7H2O c. CaCl2 6H2O d. CuCl2 7H2O Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds e. CoCl2 6H2O ANSWER: e 22. What is the name of the hydrate Ca(NO3)2 4H2O? a. calcium dinitrate tetrahydrate b. cadmium nitrate tetrahydrate c. calcium nitrate tetrahydrate d. calcium nitrate pentahydrate e. cadmium nitrate hexahydrate ANSWER: c 23. What is the formula for calcium sulfate? a. CaSO4 b. CaSO3 c. CaSO2 d. CaS e. CaS2 ANSWER: a 24. Two electrons are put close very to each other. The force they experience is a. attractive and weak. b. attractive and strong. c. repulsive and weak. d. repulsive and strong. e. nothing.

ANSWER: d 25. To form a neutral solid the positive charge must be _____________ the negative charge. a. less than b. greater than c. equal to d. far away from e. none of the above ANSWER: c 26. What is the name of the ionic compound Mn(ClO4)2? a. manganese chlorate b. manganese(II) perchlorate Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds c. manganese(II) chloride d. manganese(III) hypochlorite e. manganese dichlorate ANSWER: b 27. When two oppositely charged ions come close together, they are attracted and form what kind of bond?

ANSWER: an ionic bond 28. What is the name of the ionic compound Al2O3?

ANSWER: Aluminum oxide 29. We can calculate the force of attraction or repulsion between two charged particles using what law?

ANSWER: Coulomb's law 30. Predict the empirical formula of an ionic compound made from barium and iodide ions.

ANSWER: BaI2 31. What is the formula of sodium phosphate?

ANSWER: Na3PO4 32. Which of the following compounds has at least one covalent bond? a. NaCl b. H2O c. KBr d. MgCl2 e. CaO ANSWER: b 33. What is the formula for nitrogen trifluroride? a. NF b. N3F c. NF2 d. NF3 e. N4F3 ANSWER: d 34. What is the name for the compound SF6?

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Chapter 3 From Elements to Compounds a. sulfur fluoride b. sulfur hexafluorine c. sulfur hexafluoride d. monosulfur hexafluoride e. hexasulfur fluoride ANSWER: c 35. What type of bond forms when two pairs of electrons are shared? a. single covalent bond b. double covalent bond c. triple covalent bond d. ionic bond e. metallic bond ANSWER: b 36. What is the name of the compound CBr4? a. carbon tetrabromide b. carbon bromide c. monocarbon tetrabromide d. carbide tetrabromide e. carbon tetrabromine ANSWER: a 37. What is the name of P4O10? a. phosphorus oxide b. tetraphosphorus decoxygen c. tetraphosphorus decoxide d. phosphide oxide e. phosphorus decoxide ANSWER: c 38. What is the formula of dinitrogen monoxide? a. NO b. NO2 c. N2O d. N2O2 e. N2O4 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds ANSWER: c 39. What is the formula for silicon dioxide? a. SO2 b. SiO4 c. SO4 d. Si2O4 e. SiO2 ANSWER: e 40. Which of the following elements is NOT diatomic? a. hydroge n b. bromine c. fluorine d. nitrogen e. helium ANSWER: e 41. What is the name for SiCl4? a. silicon tetrachloride b. silicon chloride c. monosilicon chloride d. monosilicon tetrachloride e. tetrachloride silicon ANSWER: a 42. What is the name of the covalent compound S2O4?

ANSWER: disulfur tetroxide 43. What is a covalent bond?

ANSWER A covalent bond is a chemical bond where two electrons are shared between : two atoms.

44. What is the formula for xenon hexafluoride?

ANSWER: XeF6 45. What is the name for N2S5?

ANSWER: dinitrogen Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds pentasulfide 46. What type of elements usually form covalent bonds?

ANSWER: nonmetal s 47. Ionic compounds form between what two types of atoms? a. two metals b. two nonmetals c. a metal and a nonmetal d. two noble gases e. a solid and a liquid ANSWER: c 48. Which of the following is NOT a characteristic property of ionic compounds? a. high melting and boiling points b. low melting and boiling points c. high solubility in water d. good electrical conductivity in solution e. usually solid at room temperature ANSWER: b 49. Which of the following is typically TRUE for covalent compounds? a. good conductors of electricity in solution b. high melting and boiling points c. highly soluble in water d. low melting and boiling points e. always solid at room temperature ANSWER: d 50. Which of the following is an example of an ionic compound? a. N2O b. KBr c. H2O d. N2 e. C3H8 ANSWER: b 51. Which of the following is an example of a covalent compound? a. LiBr Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds b. MgO c. H2O d. MgCl2 e. NaF ANSWER: c 52. Which of the following compounds is an acid? a. H2 b. CaCl2 c. K2SO4 d. NaNO3 e. H2CO3 ANSWER: e 53. What is the name of the acid H2SO3? a. sulfurous acid b. sulfuric acid c. hydrosulfuric acid d. dihydrogen sulfite e. hydrosulfurous acid ANSWER: a 54. What is the formula for perchloric acid? a. HCl b. HClO c. HClO2 d. HClO3 e. HClO4 ANSWER: e 55. An aqueous solution of an ionic compound is called a. a heterogeneous solution. b. a magnetic solution. c. an electrochemical solution. d. an electrolyte solution. e. a conducting solution.

ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds 56. What does phosphoric acid ionize into when dissolved into water?

ANSWER: It will ionize into three hydrogen ions and one phosphate ion.

57. Based on its formula, would you expect BaCO3 to be ionic or covalent?

ANSWER Barium is a metal and the carbonate anion is made of nonmetals, which makes : barium carbonate an ionic compound.

58. How do the melting points of ionic and covalent compounds differ?

ANSWER Ionic compounds tend to have higher melting points, and covalent compounds : tend to have lower melting points.

59. How do the behaviors of ionic and covalent compounds differ in aqueous solution?

ANSWER Soluble ionic compounds dissociate into their constituent ions, whereas : covalent compounds remain as discrete molecules.

60. Would you expect an aqueous solution of copper(II) chloride to conduct electricity?

ANSWER Yes, a solution of copper(II) chloride will conduct electricity as ionic : compounds dissociate into their constituent ions, becoming strong electrolytes.

61. What is the purpose of elemental analysis? a. to determine the molecular formula b. to measure the molecular mass of a compound c. to determine the percent composition of elements in a compound d. to identify the empirical formula only e. to separate elements from a mixture ANSWER: c 62. What does the empirical formula of a compound represent? a. the actual number of atoms in a molecule b. the simplest whole-number ratio of elements in a compound c. the molecular weight of a compound d. the ratio of isotopes in a molecule e. the arrangement of atoms in a molecule ANSWER: b 63. If a compound consists of 40% carbon, 6.7% hydrogen and 53.3% oxygen by mass, what is its empirical formula? a. CHO b. CH2O c. CH3O d. C3H6O2 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 3 From Elements to Compounds e. C4H8O ANSWER: b 64. What is the molecular formula of a compound if its empirical formula is CH2 and the molar mass is 56 g/mol? a. CH2 b. C3H6 c. C2H4 d. C4H8 e. C5H10 ANSWER: d 65. Mass spectrometry provides what information about a compound? a. the percent composition b. the empirical formula c. the empirical mass d. the molecular mass e. the atomic mass ANSWER: d 66. Which of the following best describes combustion analysis? a. analysis of products formed by burning a sample b. the identification of elements through heating c. analysis of metallic elements d. separation of mixtures based on combustion products e. measurement of bond energies ANSWER: a 67. In combustion analysis, which of the following is measured to determine the amount of hydrogen in a compound? a. carbon dioxide (CO2) b. hydrogen (H2) c. oxygen (O2) d. methane (CH4) e. water (H2O) ANSWER: e 68. If a compound contains 82.6% carbon and 17.4% hydrogen by mass, what is its empirical formula?

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Chapter 3 From Elements to Compounds a. CH3 b. C2H5 c. C4H5 d. C4H12 e. C5H8 ANSWER: b 69. Which of the following is TRUE about molecular formulas? a. They always have a greater number of atoms than empirical formulas. b. They represent the simplest ratio of atoms in a compound. c. They provide the exact number of each type of atom in a molecule. d. They are the same as a structural formula. e. They are always multiples of isotopic masses.

ANSWER: c 70. What is the empirical formula of C8H8O3? a. C4H4O3 b. C8H8O3 c. C16H16O6 d. C4H4O3 e. None of the above ANSWER: b 71. A compound has an empirical formula of CH2 and a molecular mass of 97.8 g/mol, what is the molecular formula of this compound?

ANSWER The mass of the CH2 unit is 14.03 g/mol. Taking 97.8/14.03 gives a value of : 6.97 ≈ 7, so the molecular formula would be C7H14.

72. What are the three components of a mass spectrometer?

ANSWER The three main components of a mass spectrometer are: an ionizing chamber, : an electric field, and a detector.

73. A compound was analyzed and determined to be 53.3% C, 11.2% H, and 35.5% O by mass. Further analysis by mass spectrometry determined the molecular weight to be 90.1 g/mol. What is the molecular formula for this compound?

ANSWER: C4H10O2 74. An unknown compound was analyzed by combustion analysis and found to be 37.23% C, 7.81% H, and 54.96% Cl. What is the empirical formula of this compound?

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Chapter 3 From Elements to Compounds ANSWER: C2H5Cl 75. What information is required to find the molecular formula of a compound?

ANSWER: We need the empirical formula and the molecular mass to find the molecular formula.

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Chapter 2 Atoms 1. Matter is neither created nor destroyed in chemical reactions. This is a statement of a. the first law of thermodynamics. b. the law of conservation of mass. c. the law of constant composition. d. the law of multiple proportions. e. the law of conservation of energy.

ANSWER: b 2. Which of the following statements is NOT part of Dalton's atomic theory? a. Elements are made of atoms. b. The atoms of each element are unique. c. Atoms can join together in whole number ratios. d. Atoms are made up of subatomic particles. e. Atoms are unchanged in chemical reactions.

ANSWER: d 3. The law of multiple proportions states that a. atoms may be combined in ANY proportions to form different compounds. b.when two elements are able to form multiple compounds, the elements combine in different whole-number ratios. c. two elements will always form compounds with the same whole-number ratio. d.the total mass of the products is equal to the total mass of the reactants. e. atoms of the same element are identical in size and mass.

ANSWER: b 4. Over time, 15 kg of iron reacts with oxygen in the presence of water to form iron(III) oxide, or rust. If the mass of the rust is 21.45 kg, how much oxygen had to react to produce that amount of rust?

ANSWER: 6.45 kg 5. Which of the following best illustrates the law of conservation of mass? a. A cooked egg weighs less than a raw egg. b. A piece of paper burns and produces less mass in ashes. c. A cube of ice weighs the same as the puddle of water it forms after melting. d. A plant grows and increases its mass by absorbing sunlight. e. Water evaporates as its mass disappears.

ANSWER: c 6. Which of the following observations would violate the law of conservation of mass? a. Atoms rearranged during a chemical reaction. b. A sealed reaction vessel shows no change before and after a chemical reaction.

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Chapter 2 Atoms c. A burning log loses mass because gases escape into the air. d. The mass of the reactants is equal to the mass of the products for a chemical reaction. e. A sealed reaction vessel shows a decrease in mass after a chemical reaction.

ANSWER: e 7. When a chemical reaction occurs, which of the following must be TRUE about the products? a. The products will be the same as the reactants. b.The total mass of the products will be the same as the total mass of the reactants, but they don't have to be the same atoms. c. The number of atoms for each element will be the same, but they will be arranged differently. d.The number of atoms for each element will be the same, but they will be arranged in the same way. e. The number of atoms for each element cannot be determined from this information.

ANSWER: c 8. In what ratios can atoms combine to form compounds?

ANSWER: Atoms can only combine in whole number ratios.

9. If 33 g of substance A reacts with 12 g of substance B, what will the mass of the products be? a. 33 g b. 12 g c. 21 g d. 45 g e. 57 g ANSWER: d 10. The colors in fireworks come from chemicals that burn different colors. For example, red fireworks usually get their color from a strontium compound like strontium chloride (SrCl2). If 39.85 g of SrCl2 reacts with 8.11 g of O2, what will be the total mass of the products?

ANSWER The mass of the products will be the same as the reactants according to the law : of conservation of mass. Therefore, the total mass of the products is 47.96 g.

11. Based on its location in the periodic table, which of these elements is a metal? a. carbon b. helium c. nitrogen d. vanadiu m Copyright Macmillan Learning. Powered by Cognero.

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Chapter 2 Atoms e. iodine ANSWER: d 12. Which of the following elements has properties the most similar to calcium? a. potassium b. magnesiu m c. sodium d. scandium e. krypton ANSWER: b 13. Which of the following elements is a metalloid? a. aluminum b. tin c. phosphorus d. arsenic e. gallium ANSWER: d 14. Who was the person that organized the first modern periodic table?

ANSWER: Dimitri Mendeleev 15. The halogens (group 7A) tend to form what kind of molecules in their elemental form?

ANSWER: diatomic molecules 16. What is the chemical symbol for tin? a. Ti b. Sn c. Tn d. T e. Sb ANSWER: b 17. Group 8A elements are often called noble gases. What is a common property of elements in this group? a. They are inert, and rarely form compounds. b. They readily form compounds with other elements. c. They are highly reactive. d. They are good conductors of electricity. e. They are magnetic.

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Chapter 2 Atoms ANSWER: a 18. Where are nonmetals most likely to be found on the periodic table?

ANSWER: Nonmetals are most likely to be found in the upper right corner of the periodic table.

19. The inner transition metals consist of two rows of elements which are usually shown at the bottom of the periodic table. The second row of these elements is called a. rare earth metals. b. lanthanide series. c. actinide series. d. alkali earth metals. e. noble series.

ANSWER: c 20. Which of the following elements is in the lanthanide series? a. barium b. radium c. thorium d. cerium e. lead ANSWER: d 21. What is a common property of alkali metals?

ANSWER Alkali metals are in the far left of the periodic table. The alkali metals are soft : and react violently with water or air.

22. Which period of the periodic table contains the element chromium? a. period 1 b. period 2 c. period 3 d. period 4 e. period 5 ANSWER: d 23. Which of the following particles is/are found outside of the nucleus? a. protons Copyright Macmillan Learning. Powered by Cognero.

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Chapter 2 Atoms b. neutrons c. electrons d. protons and neutrons e. electrons and neutrons ANSWER: c 24. Which experiment led to the discovery of the electron? a. gold foil experiment b. oil drop experiment c. double slit experiment d. cathode ray tube experiment e. Rutherford scattering experiment ANSWER: d 25. What did Rutherford's gold foil experiment demonstrate? a. The nucleus is small, dense, and positively charged. b. The nucleus contains two subatomic particles: protons and neutrons. c. Electrons are located in specific energy levels. d. Protons are positively charged. e. Electrons have a negative charge.

ANSWER: a 26. What was the main conclusion from Millikan's oil drop experiment? a. The atom is mostly empty space. b. The charge on an electron. c. The nucleus is positively charged. d. Electrons orbit the nucleus at fixed distances around the nucleus. e. Atoms contain neutrons.

ANSWER: b 27. Which of the following subatomic particles has/have no charge? a. protons b. neutrons c. electrons d. protons and electrons e. protons and neutrons ANSWER: b 28. What is the key idea in Bohr's model of the atom?

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Chapter 2 Atoms ANSWER: Electrons orbit the nucleus like planets orbit the sun.

29. What model of the atom did J.J. Thomson propose, and how did it describe the structure of the atom?

ANSWER Thomson proposed the plum pudding model, in which atoms are solid with : negative electrons spread throughout a positively charged matrix.

30. How did Rutherford's model of the atom differ from Thomson's model?

ANSWER Rutherford's model proposed that the atom has a small, dense, positively : charged nucleus with electrons orbiting around it. While Thomson's model described the atom as a spherical solid with negative electrons spread throughout a positively charged matrix.

31. Two spheres each having a charge of +2.7 nanocoulombs are separated by a distance of 3.7 mm. Using Coulomb's law, calculate the force of repulsion between these two spheres.

ANSWER: 4.8 mN or 4.8×10–3 N 32. What was the quantity determined by the cathode-ray tube experiment? a. The mass of the electron. b. The charge of the electron. c. The mass to charge ratio of the electron. d. The charge of the proton. e. The mass to charge ratio of the hydrogen nucleus.

ANSWER: c 33. As a result of Volta's battery, scientist found that they could separate some compounds into their elements.

What did this discovery allow them to do? a. develop the periodic table b. discover new subatomic particles c. propose the law of conservation of mass d. identify how charge particles interact with each other e. invent new types of batteries ANSWER: d 34. What is the more formal name for the type of battery that Volta made? a. dry cell battery b. nickel cadmium battery c. lithium ion battery d. mercury battery e. electrochemical cell ANSWER: e 35. The atomic number of an element is defined by the Copyright Macmillan Learning. Powered by Cognero.

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Chapter 2 Atoms a. number of protons. b. number of neutrons. c. number of electrons. d. sum of the number of protons and neutrons. e. atomic mass.

ANSWER: a 36. What does the mass number of an atom represent? a. number of protons b. number of electrons c. number of neutrons d. the sum of the number of protons and neutrons e. the number of isotopes of an element ANSWER: d 37. Using the periodic table, determine how many protons and neutrons are in a bromine-79 atom.

ANSWER Bromine has an atomic number of 35 so it has 35 protons. Subtracting 35 from : 79, gives 44 neutrons.

38. Chlorine has two main isotopes: 35Cl and 37Cl. 35Cl has a mass 34.9 u, and an abundance of 75.8%. 37Cl has a mass of about 36.9 u with an abundance of about 24.2%. From this data, calculate the atomic mass of chlorine.

ANSWER: 35.38 u 39. How many neutrons are in an atom of cobalt-60? a. 27 b. 33 c. 60 d. 32 e. 87 ANSWER: b 40. What is the mass number of an atom that has 20 protons, 22 neutrons, and 20 electrons? a. 20 b. 22 c. 40 d. 42 e. 62 ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 2 Atoms 41. What is the charge of the nucleus in an atom? a. positive b. negative c. zero d. positive or negative, depending on the element e. switches between positive and negative ANSWER: a 42. What are isotopes? a. Atoms of the same element with the same number of neutrons, but a different number of protons. b. Atoms of the same element with the same number of protons, but a different number of neutrons. c. Atoms of the same element with the same number of protons, but a different number of electrons. d. Atoms of the same element with the same number of electrons, but different number of protons. e. Atoms of the same element with the same number of neutrons, but different number of electrons.

ANSWER: b 43. What element has an atomic number of 19? a. calcium b. sodium c. potassium d. fluorine e. argon ANSWER: c 44. What is the mass number of an isotope of gallium that has 40 neutrons?

ANSWER: 71 45. What is the chemical symbol for the element with the atomic number of 37? a. K b. Sr c. Cl d. Ar e. Rb ANSWER: e Copyright Macmillan Learning. Powered by Cognero.

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Chapter 2 Atoms 46. There are two naturally occurring radioactive isotopes of tellurium. An atom of radioactive tellurium has a mass number of 128, how many neutrons does it have?

ANSWER: There are 76 neutrons in this atom.

47. Dichloromethane is a common solvent used in organic chemistry. What is the formula mass of dichloromethane (CH2Cl2)? a. 48.96 u b. 62.93 u c. 84.93 u d. 96.00 u e. 100.50 u ANSWER: c 48. What is the percent composition of hydrogen in acetic acid (C2H4O2)? a. 3.4% b. 1.6% c. 6.7% d. 40% e. 53% ANSWER: c 49. Hydrogen peroxide has the chemical formula H2O2. If you have 2.00 kg of hydrogen peroxide, how much oxygen by mass do you have? a. 0.12 kg b. 0.24 kg c. 1.00 kg d. 1.88 kg e. 2.00 kg ANSWER: d 50. Sodium bicarbonate, NaHCO3, is commonly used to clean up acid spills in chemistry labs, and to provide leavening to some baked goods. What is the formula mass of this compound? a. 52.01 u b. 68.01 u c. 72.00 u d. 84.01 u e. 104.02 u ANSWER: d 51. What is the formula mass for zinc fluoride (ZnF2)?

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Chapter 2 Atoms ANSWER: 103.38 u 52. What is the percent mass of oxygen in calcium sulfate (CaSO4)?

ANSWER: 47.0% 53. What is the formula mass of propane (C3H8)? a. 13.02 u b. 20.08 u c. 37.04 u d. 44.11 u e. 56.14 u ANSWER: d 54. What is the formula mass of sodium sulfide, Na2S? a. 55.06 u b. 78.05 u c. 81.13 u d. 87.13 u e. 119.2 u ANSWER: b 55. What is the percent mass of chlorine in FeCl3?

ANSWER: 65.6% Cl 56. Aluminum oxide (Al2O3) forms when aluminum is exposed to oxygen. What is the mass percent of aluminum in aluminum oxide? a. 26.47% Al b. 47.38% Al c. 52.92% Al d. 62.79% Al e. 73.88% Al ANSWER: c 57. Sodium nitrate has the formula NaNO3. What is the formula mass of sodium nitrate? If you have a 5.00 kg jar of sodium nitrate in the laboratory stockroom, how much of that 5.00 kg bottle is sodium by mass?

ANSWER The formula mass of sodium nitrate is 85.00 u. Sodium makes up : approximately 1.35 kg of the 5.00 kg of sodium nitrate.

58. How many atoms are in a molecule of acetic acid (HC2H3O2)? a. 3 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 2 Atoms b. 4 c. 6 d. 7 e. 8 ANSWER: e 59. In the compound NaHSO2, which element contributes the least to the mass of the compound? a. hydrogen b. oxygen c. sodium d. sulfur e. sulfur or oxygen ANSWER: a 60. If you have one mole of carbon dioxide, how many molecules of carbon dioxide do you have? a. 1 molecule b. 44.01 molecules c. 6.02×1023 molecules d. 1.66×1024 molecules e. 2.40×1024 molecules ANSWER: c 61. How many moles contain 3.01×1023 molecules of HCl? a. 0.250 moles b. 0.500 moles c. 2.00 moles d. 8.26×1021 moles e. 1.81×1047 moles ANSWER: b 62. How many molecules are in 3.00 moles of HNO3? a. 4.98×1024 molecules b. 6.02×1023 molecules c. 1.81×1024 molecules d. 3 molecules e. 2.01×1023 molecules ANSWER: d Copyright Macmillan Learning. Powered by Cognero.

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Chapter 2 Atoms 63. What is the molar mass of CH3OH? a. 20.04 g/mol b. 29.02 g/mol c. 31.04 g/mol d. 32.05 g/mol e. 78.03 g/mol ANSWER: d 64. How many molecules are in 0.233 moles of carbon monoxide (CO)? a. 3.87×1025 molecules CO b. 2.58×1024 molecules CO c. 6.52 molecules CO d. 1.40×1023 molecules CO e. 8.32×10–3 molecules CO ANSWER: d 65. How many hydrogen atoms are in one mole of acetic acid (HC2H3O2)? a. 6.02×1023 H atoms b. 1.51×1023 H atoms c. 2.01×1023 H atoms d. 1.81×1024 H atoms e. 2.40×1024 H atoms ANSWER: e 66. If a sample of calcium iodide (CaI2) has 7.98×1024 formula units, what is the mass of calcium iodide in the sample? a. 0.0132 kg b. 0.787 kg c. 1.63 kg d. 2.21 kg e. 3.90 kg ANSWER: e 67. Phosphoric acid has the formula H3PO4. If you have one mole of phosphoric acid, how many moles of atoms do you have? a. 1 b. 3 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 2 Atoms c. 4 d. 7 e. 8 ANSWER: e 68. How many moles contain 5.28×1023 atoms of copper?

ANSWER: 0.877 moles Cu 69. What is the molar mass of phosphoric acid (H3PO4)?

ANSWER: 98.00 g/mol 70. If you have 873.22 g of aluminum oxide, how many moles of aluminum oxide do you have? The molar mass aluminum oxide is 101.96 g/mol.

ANSWER: 8.56 moles Al2O3 71. Potassium iodide has a molar mass of 166.00 g/mol. If you have 0.611 moles of potassium iodide, how many grams of potassium iodide do you have?

ANSWER: 101.43 g KI 72. How many molecules are in 5.00 moles of O2? a. 1.20×1024 molecules b. 8.31×10–24 molecules c. 3.01×1024 molecules d. 6.02×1024 molecules e. 9.03×1025 molecules ANSWER: c 73. What is the mass of 7.32×1024 formula units of sodium sulfate (Na2SO4)? a. 1720 g b. 12.2 g c. 863 g d. 9.79 g e. 5.84 g ANSWER: a 74. How many formula units are in a 35.00 g sample of magnesium sulfate (MgSO4)?

ANSWER: 1.75×1023 formula units of MgSO4 Copyright Macmillan Learning. Powered by Cognero.

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Chapter 2 Atoms 75. A given sample of iron(II) bromide (FeBr2) contains 1.99×1022 formula units of iron(II) bromide. How many grams of iron(II) bromide is this?

ANSWER: 7.12 g FeBr2

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Chapter 1 Matter and Measurement 1. Which of the following career fields are connected to chemistry? a. biology and medicine b. physics and engineering c. agriculture and food science d. environmental, earth, and space science e. all of the above ANSWER: e 2. A scientific law describes an observation is widely considered true. What else is TRUE of a scientific law? a. It will tell us why the observation is true. b. It can be used to make predictions. c. It is a theory that has been proven true. d. All of the above are true. e. None of the above are true.

ANSWER: b 3. Which of the following is NOT true about a theory? a. A theory can be considered a way of thinking about a particular topic. b. A theory is an idea that is supported by experimental evidence. c. A theory is part of the scientific method. d. A theory can change based on new evidence. e. A theory is an idea that is completely unproven.

ANSWER: e 4. A hypothesis is a. a proven fact. b. an experimental procedure. c. a tentative explanation that has not been tested. d. an idea that has been tested and refined. e. a statement that describes observations that are true in widely varying circumstances.

ANSWER: c 5. According to the scientific method, what is the next step after developing a hypothesis? a. draw conclusions b. devise and conduct experiments c. develop a theory d. share results with other scientists e. develop another hypothesis Copyright Macmillan Learning. Powered by Cognero.

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Chapter 1 Matter and Measurement ANSWER: b 6. Which of the following is TRUE about the scientific method? a. First, observations are made, then new ideas are formulated, and it ends with testing. b. The first step is to make conclusions. c. It is a cyclic process. d. The last step is to run experiments. e. None of the above.

ANSWER: c 7. Chemistry is often called the central science. Why? ANSWER Chemistry is used in many other scientific fields such as biology, physics, : agriculture, and engineering.

8. What is chemistry?

ANSWER: Chemistry is the study of what substances are made of and how they behave.

9. How do you interact with chemistry in your daily life?

ANSWER: Chemistry is everywhere. Everything you do is impacted by chemistry!

10. Briefly explain the scientific method.

ANSWER The scientific method is a cyclic process of making observations, formulating : new ideas, and then testing those ideas with experiments.

11. An element is defined as a. a substance that is made of only one type of atom. b. anything that has mass and takes up volume. c. a substance that can be broken down into simpler substances. d. a combination of two or more pure substances. e. a substance composed of two or more types of atoms bound together in a fixed ratio.

ANSWER: a 12. A compound is a substance that a. contains more than one type of atom present in fixed ratios. b. contains only one type of atom. c. contains more than one type of atom present in different ratios. d. can contain one or many types of atoms. e. must contain at least three different types of atoms.

ANSWER: a Copyright Macmillan Learning. Powered by Cognero.

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Chapter 1 Matter and Measurement 13. Which of the following is an element? a. water (H2O) b. oxygen (O2) c. sucrose (C12H22O11) d. sodium chloride (NaCl) e. carbon dioxide (CO2) ANSWER: b 14. Which of the following is a compound? a. copper (Cu) b. fluorine (F2) c. neon (Ne) d. carbon (C) e. nitrogen monoxide (NO) ANSWER: e 15. Which of the following is an example of a heterogeneous mixture? a. salt water b. sugar dissolved in coffee c. air d. salad e. ketchup ANSWER: d 16. Which of the following characteristics best describes a liquid? a. Has a definite volume and a definite shape. b. Does not have a definite volume nor shape. c. Has a definite volume, but not a definite shape. d. The particles vibrate within their spaces in a framework. e. The particles move freely with very little interaction with each other.

ANSWER: c 17. What is the process of a liquid transitioning into a gas called? a. melting b. vaporization c. freezing d. sublimation e. condensation ANSWER: b Copyright Macmillan Learning. Powered by Cognero.

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Chapter 1 Matter and Measurement 18. Which of the following is an example of a homogenous mixture? a. a soda with ice b. oil and water c. chicken noodle soup d. saltwater e. sand and rocks ANSWER: d 19. Which of the following is an example of a chemical change? a. You boil some water for pasta. b. You pour the excess pasta water down the drain. c. The steam from the water turns to liquid on your glasses. d. You wipe the water from your glasses. e. You digest your food.

ANSWER: e 20. If liquid water is frozen into solid ice, what is happening on a molecular level? a. The molecules get closer together and have less energy. b. The molecules get further apart and have less energy. c. The molecules get closer together and have more energy. d. The molecules get further apart and have more energy. e. No change in distance or energy occurs.

ANSWER: b 21. What happens to the starting substances as they are converted into products in an endothermic reaction? a. The energy of the substances does not change. b. The energy of the substances decreases. c. The energy of the substances increases. d. The substances become less reactive. e. The substances become more stable.

ANSWER: c 22. What type of mixture has a composition that is the same throughout?

ANSWER: Homogenous mixture 23. What is the difference between the composition and structure of a material?

ANSWER The composition of a material is the components that make up the material. : The structure of a material is how those components are arranged.

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Chapter 1 Matter and Measurement 24. What are the differences between atoms, elements, and compounds?

ANSWER Atoms are the fundamental unit of matter. All matter is made up of atoms. An : element is the simplest form of matter. Every element is made up of one type of atom. Compounds are composed of more than one type of atom, bound together in fixed ratios.

25. If a chemical change has occurred, what must be true?

ANSWER If a chemical change has occurred, a chemical reaction has occurred, and we : have made a new, chemically distinct substance.

26. Water is pumped to the top of a water tower to increase the municipal water pressure. Does the water at the top of the tower have more or less potential energy than the water at the base of the tower?

ANSWER The water at the top has more potential energy than the water at the base due to : its position.

27. Which of the following is NOT an SI base unit? a. meters b. seconds c. amperes d. grams e. candela ANSWER: d 28. What is the SI unit of temperature? a. degree Celsius b. degree Fahrenheit c. Kelvin d. ampere e. joule ANSWER: c 29. What is the prefix for

in the SI system?

a. centib. kiloc. microd. millie. megaANSWER: d

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Chapter 1 Matter and Measurement 30. The typical human body temperature is 98.6 F. What is this temperature in Kelvin? a. 310.2 K b. 37.0 K c. 371.8 K d. 209.5 K e. 482.6 K ANSWER: a 31. Which of the following is the SI unit for length? a. centimeter b. foot c. meter d. kilometer e. mile ANSWER: c 32. Convert 315.15 K to degrees Celsius. a. 588.30 C b. 107.6 C c. 42.00 C d. 5.556 C e. 315.15 C ANSWER: c 33. How many megagrams are in 1 gram? a.

Mg b.

Mg

c.

Mg

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Chapter 1 Matter and Measurement d.

Mg

e.

Mg

ANSWER: c 34. What is the volume of a sphere with a radius of 3.2 cm? The volume of a sphere is given by

.

a. 44 cm3 b. 140 cm3 c. 43 cm3 d.

cm3

e. 78 cm3 ANSWER: b 35. How many milligrams are equal to 1 gram? a. 0.001 mg b. 0.01 mg c. 100 mg d. 1000 mg e. 1,000,000 mg ANSWER: d 36. Which of the following substances will sink in water (d = 1.0 g/mL)? a. acetone (d = 0.78 g/mL) b. vegetable oil (d = 0.92 g/mL) c. honey (d = 1.42 g/mL) d. All of the above will sink in water. e. None of the above will sink in water.

ANSWER: c 37. How many micrometers are in 1 meter?

ANSWER: There are 1,000,000 micrometers in 1 meter.

38. Why is it important to use correct units consistently?

ANSWER Measurements don't have any meaning without a unit attached. Units of : measurement are quantities with accepted values that can be used to communicate between people. Additionally, units can help you check whether Copyright Macmillan Learning. Powered by Cognero.

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Chapter 1 Matter and Measurement you've done a calculation correctly.

39. The temperature of a typical refrigerator is about 4 C. What is this temperature in degrees Fahrenheit?

ANSWER:

40. A cube has a side length of 5.3 cm and a mass of 125.8 g. Will this cube float in water? The density of water is 1.0 g/cm3.

ANSWER: d = 0.84 g/cm3, so the cube will float in water 41. A cube with a side length of 5.00 cm has a mass of 131 g. Calculate the volume (in cm3) and density (in g/cm3) of the cube.

ANSWER:

42. A student measures 210.50 mL of water for an experiment in her chemistry lab. How many significant digits does 210.50 have? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: e 43. Which of the following measurements has the highest precision? a. 15.0 cm b. 15.000 cm c. 15 cm d. 15.00 cm e. 1.5 cm ANSWER: b 44. If a set of measurements are all very close to each other, but are far away from the true value, they are a. accurate and precise. b. accurate but not precise. c. precise but not accurate. d. neither accurate nor precise.

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Chapter 1 Matter and Measurement e. none of the above.

ANSWER: c 45. Which of the following numbers has three significant digits? a. 0.12 b. 3.850 c. 305 d. 432.0 e. 0.001 ANSWER: c 46. When multiplying 3.45 by 2.9, how many significant digits should the result have? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: b 47. The height of a book is measured as 8.0230 cm. How many significant digits are in this measurement? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: e 48. How many significant digits are in the measurement 0.0930? a. 1 b. 2 c. 3 d. 4 e. 5 ANSWER: c 49. The mass of molten aluminum that is needed to make a specific car part is 275 g. Which of the following measurements is the most accurate? a. 280 g b. 278.714 g c. 288.32 g Copyright Macmillan Learning. Powered by Cognero.

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Chapter 1 Matter and Measurement d. 300 g e. 276.4 g ANSWER: e 50. A pallet of lumber has 256 boards on it. If the mass of the boards on the pallet weighs 928.96 kg, how much does each board weigh using an appropriate number of significant digits? a. 3.6288 kg b. 3.629 kg c. 3.63 kg d. 3.6 kg e. 4 kg ANSWER: a 51. What kind of numbers are exempt from significant digits? a. measured numbers b. decimal numbers c. exact numbers d. large numbers e. numbers in scientific notation ANSWER: c 52. Kris walks 1.32 km from his home to the hardware store, 1.1 km from the hardware store to the grocery store, and 1.435 km from the grocery store back to home. What is the total distance he traveled on this trip using appropriate significant digits? a. 3.86 km b. 3.9 km c. 4 km d. 4.0 km e. 3.855 km ANSWER: b 53. Accuracy and precision are terms that scientists use to describe the qualities of a data set. What does the accuracy of a data set indicate?

ANSWER: The accuracy of a data set indicates how well the data reflects the true value.

54. Calculate the area of a rectangle with side lengths of 3.042 cm and 2.1 cm. Report your answer with appropriate units and significant digits.

ANSWER: 55. A student adds 5.0 mL of water to a graduated cylinder. They then add an insoluble solid to the graduated cylinder and record the volume of the water and solid as 9.15 mL. What is the volume of the solid reported Copyright Macmillan Learning. Powered by Cognero.

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Chapter 1 Matter and Measurement using appropriate significant figures?

ANSWER: 56. If a set of measurements reflects the true value of a measurement well, but the individual measurements are scattered away from the true value, what can be said about the accuracy and precision of these measurements?

ANSWER The set of measurements is accurate (reflects the true value), but the : measurements are not precise (they are far away from each other).

57. If you have a pencil that weighs 4.632 g, a pen that weighs 7.15 g, and a stylus that weighs 20.7 g, what is the total mass of these three objects? Report the answer using correct significant digits.

ANSWER: 58. Ibuprofen is a commonly used over-the-counter medication used to treat body aches and pains. A common dose of ibuprofen is 400 mg. What is the mass of this dose of ibuprofen in grams? a. 40 g b. 4 g c. 0.4 g d. 0.04 g e. 0.004 g ANSWER: c 59. Convert 100 m to kilometers (km). a. 0.001 km b. 0.1 km c. 1 km d. 10,000 km e. 100,000 km ANSWER: b 60. How many centimeters are in 3.4 m? a. 0.0034 cm b. 0.034 cm c. 340 cm d. 3400 cm e. 34000 cm ANSWER: c 61. Convert 4.7 L to milliliters. a. 0.0047 mL b. 0.047 mL Copyright Macmillan Learning. Powered by Cognero.

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Chapter 1 Matter and Measurement c. 4700 mL d. 47000 mL e. 470000 mL ANSWER: c 62. The speed of light in a vacuum is a.

m/s. What is this value in miles per hour?

miles per hour

b. 51.6 miles per hour c. miles per hour d.

miles per hour

e.

miles per

hour ANSWER: d 63. An Olympic-size swimming pool is 50. m × 25 m. What is the surface area of the pool in square feet? a. 4100 ft2 b. 1300 ft2 c. 120 ft2 d. 13,000 ft2 e. 380 ft2 ANSWER: d 64. Which of the following is the conversion of 100 mL to cm3? a. 1 cm3 b. 10 cm3 c. 100 cm3 d. 1000 cm3 e. 10000 cm3 ANSWER: c 65. How many liters are in 1 gallon of lemonade? 1 L = 0.264 gal Copyright Macmillan Learning. Powered by Cognero.

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Chapter 1 Matter and Measurement a. 0.264 L b. 1.26 L c. 2.32 L d. 3.79 L e. 7.58 L ANSWER: d 66. Water is one of the few substances whose solid phase is less dense than its liquid phase. The density of liquid water is about 1.00 g/mL, whereas the density of ice is about 0.920 g/mL. Approximately how much does a 100.00 mL sample of ice weigh? a. 100. g b. 92.0 g c. 109 g d. 100.92 g e. 99.08 g ANSWER: b 67. How many square meters, m2, are in 150 ft2? a. 46 m2 b. 14 m2 c. 490 m2 d. 1600 m2 e. 0.30 m2 ANSWER: b 68. A water tank has a volume of 5 m3. What is this volume in liters? a. 0.005 L b. 0.05 L c. 50 L d. 500 L e. 5000 L ANSWER: e 69. A bag of pretzels has 110 kilocalories per serving. How many kilojoules are in one serving? 4.184 kJ = 1 kcal a. 460 kJ b. 26 kJ c. 156 kJ d. 440 kJ Copyright Macmillan Learning. Powered by Cognero.

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Chapter 1 Matter and Measurement e. 4.2 kJ ANSWER: a 70. What is the mass of an iron block that has a volume of 24 cm3? The density of iron is 7.87 g/cm3.

ANSWER: 190 g 71. Convert 142 mg to kilograms (kg).

ANSWER: kg

72. A student is trying to convert a measurement in millimeters to meters, but after performing the calculation, gets units of mm2/m. What mistake was made? How would the student correct this?

ANSWER The reciprocal of the conversion factor was used. The conversion factor should : be flipped.

73. The flow of coolant through a pipe is 4.50 mL/s. What would the flow rate be in liters per hour?

ANSWER: 16.2 L/h 74. A cylindrical gasoline storage tank has a radius of 1.5 m, and a height of 3.0 m. What is the volume of this tank in liters? The volume of a cylinder is given by .

ANSWER: 21,000 L 75. The density of a certain motor oil is 0.860 g/cm3. How much would 175 L of this oil weigh (in kg)?

ANSWER: 151 kg

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