TEST BANK for Quantitative Chemical Analysis 10th Edition by Daniel C. Harris and Charles A. Lucy.

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Chap 01_10e Indicate the answer choice that best completes the statement or answers the question.

1. Which constant is NOT used to define the fundamental SI units? a. Avogadro’s number b. speed of light in vacuum c. elementary charge d. Planck’s constant e. π 2. A satellite in low Earth orbit with a circular orbit has an orbital speed of 7.3 km/s relative to the Earth’s surface. Calculate the satellite’s speed in miles per hour. (1 mi = 1.609 km) a. 1.6 × 104 mi/h b. 1.3 × 10−3 mi/h c. 4.2 × 104 mi/h d. 3.3 × 10−3 mi/h e. 3.1 × 102 mi/h 3. The calorie content of a candy bar is 230. Calories per serving (1 bar). Calculate the specific energy (kJ/g) of the candy bar. (1 candy bar = 52.7 g, 1 Calorie = 1 000 calories, 1 calorie = 4.184 J) a. 2.90 × 103 kJ/g b. 18.3 kJ/g c. 1.83 × 10−2 kJ/g d. 5.07 × 104 kJ/g e. 9.59 × 10−1 kJ/g

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Chap 01_10e 4. Arrange the molecular views of four different solutions in order of increasing concentration. Diamond shapes represent solute particles, and circle shapes represent solvent particles.

a. I < II < III < IV b. II < III < I < IV c. IV < II < III < I d. IV < I < III < II e. III < IV < II < I

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Chap 01_10e 5. A student needs to prepare 500.0 mL of a solution containing 0.999 g of solid copper(II) sulfate. Which of the statements regarding the proper procedure to prepare this solution are FALSE? I. The 0.999 g of solid copper(II) sulfate is added to a 500.0-mL volumetric flask containing 500.0 mL of distilled water.

II. The 0.999 g of solid copper(II) sulfate is added to a 500.0-mL volumetric flask containing approximately 400 mL of distilled water before dilution to 500.0 mL. III. The 0.999 g of solid copper(II) sulfate is placed in an empty 500.0-mL volumetric flask, diluted to 500.0 mL with distilled water, and allowed to dissolve. a. I and II b. II and III c. I and III d. I, II, and III e. None of the statements is false.

6. The planet Mars orbits 2.279 1011 m from the Sun. Express the distance using the appropriate prefix. a. 227.9 Gm b. 227.9 mM c. 2.279 km d. 22.79 nm e. None of these is correct.

7. Which of the following is NOT a fundamental SI unit of a quantity? a. second (s) b. meter (m) c. gram (g) d. ampere (A) e. mole (mol) 8. The gas mileage for a new car model destined for sale in Europe must be determined for regulatory and promotional purposes. If the car uses 10.5 gallons to travel 250. miles, what is the gas mileage in km/L? (1 mi = 1.609 km, 1 gal = 3.785 L) a. 174 km/L b. 3.25 km/L c. 67.3 km/L d. 14.9 km/L e. 10.1 km/L

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Chap 01_10e 9. Calculate the mass of Na2CO3 (FM 105.988 8) needed to prepare a 15.00 mM solution with a volume of 500.0 mL. a. 1.258 g b. 3.180 g c. 0.794 9 g d. 7.076 g e. 0.014 1 g 10. How many grams of CaCO3 (FM 100.086) are needed to prepare 150.0 mL of an 80.0-ppm Ca2+ solution? a. 0.012 0 g b. 0.030 0 g c. 1.875 g d. 0.533 g e. 29.9 g 11. What volume of a 36.0 wt% HCl (FM 36.458) solution must be diluted to prepare 1.000 L of a 0.100 0 M HCl solution? The density of 36.0 wt% HCl is 1.18 g/mL. a. 11.7 mL b. 8.58 mL c. 1.20 mL d. 10.1 mL e. 64.6 mL 12. A mixture of 50.00 g propane (C3H8, FM 44.10) and 100.00 g oxygen (O2, FM 31.998) is combusted to form carbon dioxide and water. _________ is the limiting reactant, and ________ of _________is in excess.

C3H8(g) + 5O2(g) → 3CO2 (g) + 4H2O (g) a. Propane; 22.44 g; oxygen b. Oxygen; 63.71 g; propane c. Oxygen; 22.42 g; propane d. Propane; 27.44 g; oxygen e. None of these answers is correct.

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Chap 01_10e 13. Calculate the mass of NaCH3CO2 contained in 500.0 mL of a 0.150 0 M NaCH3CO2 solution.

(NaCH3CO2 = 82.034 3 g/mol) a. 914.3 µg b. 283.4 g c. 24.61 µg d. 6.153 g e. 24.61 g 14. Calculate the molarity of a 2.0-ppm Mg2+ solution. a. 8.2 × 10−5 M b. 8.2 × 10−2 M c. 1.2 × 10−2 M d. 1.2 × 10−5 M e. 4.9 × 10−2 M 15. The sulfur content of an ore is determined gravimetrically by reacting the ore with concentrated nitric acid and potassium chlorate, which converts all of the sulfur to sulfate. The excess nitrate and chlorate is removed by reaction with concentrated hydrochloric acid, and the sulfate is precipitated using Ba2+.

Ba2+(aq) + (aq) → BaSO4(s) Analysis of 10.183 0 g of a sulfur-containing ore yielded 13.022 1 g of BaSO4 (FM 233.43). What is the percent by mass sulfur in the ore? a. 32.18% b. 52.63% c. 10.74% d. 17.56% e. The answer cannot be calculated with available data.

16. Which statements are TRUE regarding the expression of the concentration of a 54.9-ppm Fe solution in terms of molarity?

I. The molar mass of iron is needed to calculate the moles of iron in solution.

II. The density of iron is needed to calculate the mass of iron in solution.

III. The solution density is needed to calculate the solution volume.

IV. The type of glassware used to prepare the solution must be known. a. I, III, and IV b. I and II c. I and III d. II and III e. None of these statements is true.

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Chap 01_10e 17. Calculate the molarity of a 30.0 wt% hydrogen peroxide (H2O2, FM 34.014 7) solution. The density of 30 wt% hydrogen peroxide is 1.135 g/cm3. a. 7.77 M b. 0.0100 M c. 0.100 M d. 10.0 M e. 8.82 M 18. When solutions of Pb2+ and

are mixed, the precipitate PbCrO4 is produced. What volume of 0.175

0 M CrO42− removes all Pb2+ from 50.00 mL of a 0.340 0 M Pb2+ solution? a. 97.14 mL b. 25.74 mL c. 48.57 mL d. 194.3 mL e. 75.00 mL 19. What volume of 12.1 M HCl must be diluted to prepare a 0.250 0 M HCl solution with a volume of 2.000 L? a. 41.3 mL b. 96.8 mL c. 10.3 mL d. 24.2 mL e. 6.05 mL 20. Which of the following statement(s) is/are TRUE regarding the properties of the limiting reagent in a chemical reaction?

I. The limiting reagent in a chemical reaction is the one that is consumed first.

II. Once the limiting reagent in a chemical reaction is gone, the reaction ceases.

III. The limiting reagent in a chemical reaction is the one that has the least mass. a. I b. II c. III d. I and II e. I, II, and III

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Chap 01_10e 21. What volume of a 50.0 wt% NaOH (FM 40.00) solution is needed to prepare a 0.350 0 M NaOH solution with a volume of 500. mL? The density of 50 wt% NaOH solution is 1.515 g/mL at 25°C.

22. What volume of a 25.0 mM Li+ solution is needed to prepare 100.0 mL of a 10.0 ppm Li+ solution?

23. An NaCl (FM 58.44) solution has a concentration of 33.5 wt% and a density of 1.049 2 g/mL. What is the molarity of the solution?

24. Lead(II) carbonate precipitates when aqueous lead(II) is mixed with aqueous carbonate.

Pb2+(aq) + (aq) ® PbCO3(s) If 5.000 g Pb(NO3)2 (FM 331.2) and 2.500 g Na2CO3 (FM 105.988 8) are mixed in water, which ion is the limiting reactant? What mass of PbCO3 (FM 267.21) is precipitated?

25. The recommended daily allowance of calcium for men between the ages of 19 and 50 is 1000 mg Ca. Three multivitamin tablets are analyzed for calcium gravimetrically with the precipitation of Ca2+ by the oxalate ion, . If the mass dry calcium oxalate (FM 128.097) obtained is 2.013 6 g, how many tablets must a man take in a given day to meet the recommended daily allowance?

Ca2+(aq) + (aq) → CaC2O4(s)

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Chap 01_10e 26. A 15.3-g sample of an organic compound is completely combusted in air, producing 21.0 g CO2 and 8.61 g H2O. What is the weight percent of C in the organic compound?

27. Tidal volume is the amount of air breathed in with each normal breath. The average tidal volume is 0.50 L, and the average breathing rate is 12 breaths/min. Calculate the total volume (in m3) of air a person breathes in one hour.

28. The maximum contaminant level for arsenic is 0.010 ppm for drinking water per EPA regulation. The arsenic concentration for the drinking water of a municipality was measured to be 4.92 x 10−6 M arsenic. What is the arsenic concentration of the water sample in ppm? Does the water sample meet EPA guidelines? Assume the drinking water sample has a density of 1.000 0 g/mL.

29. Find the molarity and molality of a 44.0 wt% H2SO4 (FM 98.079) solution with a density of 1.338 g/mL.

30. On average, one gallon of kerosene contains 135 000 BTU of heat energy per gallon combusted. Convert the energy content of kerosene to SI units. (1 BTU = 1 055 J, 1 gal = 3.785 L)

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Chap 01_10e Answer Key 1. e 2. a 3. b 4. d 5. c 6. a 7. c 8. e 9. c 10. b 11. b 12. c 13. d 14. a 15. d 16. c 17. d 18. a 19. a 20. d 21. 9.24 mL 22. 5.76 mL 23. 6.01 M 24. Pb2+ is the limiting reactant and 4.035 g PbCO3 is precipitated.

25. 0.210 00 g Ca per tablet; 5 vitamin tablets 26. 37.5% Copyright Macmillan Learning. Powered by Cognero.

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Chap 01_10e 27. 0.36 m3/h 28. 0.369 ppm; exceeds EPA regulation 29. 6.00 M; 6.80 m 30. 3.76 × 107 J/L or 37.6 MJ/L

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Chap 02_10e Indicate the answer choice that best completes the statement or answers the question.

1. _________________ is used to convert a precipitate to a known, constant composition. a. Ashless filter paper b. A fritted-glass funnel c. Ignition d. A rubber policeman e. A dessicator 2. _____________ is the upward force exerted on an object in a gaseous or liquid fluid. The mass measured by an analytical balance in air is _____________ its actual mass. a. Buoyancy; heavier than b. Buoyancy; lighter than c. Electromagnetic force; heavier than d. Electromagnetic force; lighter than e. Tare; equal to 3. On a lab quiz, a student listed the steps to properly use a pipet. Which step is INCORRECT? a. Use a rubber bulb to twice pull up a volume of liquid past the calibration mark and discard the contents into a waste container. b. Pull up a third volume past the calibration mark and quickly replace the bulb with the index finger. c. Touch the tip of the pipet to the side of a beaker and use the index finger to drain the liquid until the meniscus reaches the center of the calibration mark. d. Transfer the pipet to the receiving vessel, touch the tip of the pipet to the side of the vessel, and allow the pipet to drain by gravity. e. Use the rubber bulb to blow any remaining liquid from the pipet.

4. A researcher dispenses distilled deionized water from a 20-mL transfer pipet into an empty 8.437 6-g weighting bottle. If the total mass of water and weighting bottle is 28.584 5 g, what is the volume of the water delivered by the 20-mL pipet? The density of water is 0.996 786 7 g/mL. a. 20.21 mL b. 20.08 mL c. 28.68 mL d. 19.94 mL e. 19.90 mL

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Chap 02_10e 5. Which scenario has the lowest relative uncertainty? a. delivering 35.50 mL of titrant with a 50 ± 0.05 mL class A buret b. delivering 15.40 mL of titrant with a 50 ± 0.05 mL class A buret c. delivering 18.50 mL of titrant with a 25 ± 0.03 mL class A buret d. delivering 5.40 mL of titrant with a 25 ± 0.03 mL class A buret e. delivering 97.30 mL of titrant with a 100 ± 0.10 mL class A buret 6. A student titrated extracted chloride from a soil sample with 0.1 M silver nitrate. During the titration he performed the following actions while operating the buret.

Washed the buret with silver nitrate solution Drained the titrant slowly Delivered a fraction of a drop near end point Read the bottom of the concave meniscus Avoided parallax Accounted for graduation thickness in the readings His instructor notes on his lab report that the student forgot to ____________ when operating his buret. a. eliminate air bubbles b. estimate the buret reading to 1/10 of a division c. fill the buret to exactly 0.00 mL d. eliminate air bubbles and fill the buret to exactly 0.00 mL e. eliminate air bubbles and estimate the buret reading to 1/10 of a division 7. A 0.150 0 M HCl solution was prepared on a day when the temperature was 20°C. What is the concentration of the solution when used the next day at 27°C? The density of water is 0.998 207 1 g/mL at 20°C and 0.996 516 2 g/mL at 27°C. a. 0.150 3 M b. 0.149 7 M c. 0.150 8 M d. 6.653 M e. 6.632 M

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Chap 02_10e 8. Which practices reflect the safe, ethical handling of chemicals and waste in a laboratory?

I. Before working, familiarizing yourself with safety features of your laboratory II. Cleaning up spills immediately to prevent accidental contact by the next person who comes along III. Recycling of chemicals rather than disposing of waste IV. Not eating or drinking in the lab a. I, II, and IV b. II and IV c. I and II d. III and IV e. I, II, III, and IV 9. Which of the following are sources of weighing error?

I. Weighing a sample that is warmer than ambient temperature II. Cooling a sample in a desiccator prior to weighing III. Periodically calibrating the balance IV. The temperature of the balance changing over time a. I and II b. I and III c. II and IV d. II and III e. I and IV 10. ___________ is the smallest increment of mass that can be indicated by an electronic balance. a. Sensitivity b. Linearity c. Readability d. Selectivity e. Tare 11. The ____________ fulfills the critical function of reporting what a researcher has done and what she observed, and allows another researcher to repeat the work. a. lab report b. lab notebook c. MSDS d. project report e. technical note

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Chap 02_10e 12. _____________ are calibrated to deliver one fixed volume and are ____________ than _____________. a. Measuring pipets; more accurate; transfer pipets b. Transfer pipets; less accurate; measuring pipets c. Measuring pipets; less accurate; transfer pipets d. Transfer pipets; more accurate; measuring pipets e. Measuring pipets; more precise; transfer pipets 13. A student prepares a solution using a 1-L volumetric flask. When he finishes, the meniscus is above the calibration mark on the flask neck. The concentration of the solution is __________ the calculated concentration. a. less than b. greater than c. the same as d. irrelevant compared to e. impossible to compare to 14. A small air bubble trapped beneath the stopcock of a buret before a titration was expelled during the titration. Due to the air bubble, the true concentration of the solution that was titrated is ____________ the concentration calculated using the titrant volume. a. less than b. greater than c. the same as d. unrelated to e. impossible to compare to 15. Which are good practices when keeping a laboratory notebook?

I. Using complete sentences when writing notes II. Writing a balanced chemical equation for every reaction used III. Pasting hard copies of important data into the notebook IV. Recording the names of computer files where programs and data are stored a. I, II, and IV b. II and IV c. I and II d. III and IV e. I, II, III, and IV

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Chap 02_10e 16. Drying to constant mass is a common gravimetric analysis technique. Which of the following are sources of false weights?

I. A warm crucible II. Touching the crucible with bare fingers III. Using a microwave oven to dry reagents and crucibles IV. Using a desiccator a. I, II, and III b. II, III, and IV c. I and II d. II and IV e. III and IV 17. Which statement regarding volumetric flasks is FALSE? a. Volumetric flasks are calibrated to obtain a particular volume at 20°C. b. Volumetric flasks are calibrated to deliver their indicated volume. c. To properly use a volumetric flask, dissolve the reagent in less than final volume of liquid and then dilute to volume. d. The volume of the flask changes with temperature because liquid and glass expand when heated. e. To obtain the calibrated volume, the bottom of the meniscus is aligned to the center of the mark on the neck of the flask.

18. ___________ is the liquid from which a substance precipitates or crystallizes. a. Filtrate b. Eluate c. Effluent d. Mother liquor e. Slurry 19. Which statements are TRUE? I. Organic solvents, concentrated acids, and concentrated ammonia should be handled in a fume hood.

II. A respirator should be worn when handling organic solvents.

III. All containers should be labeled to indicate what they contain.

IV. Contact lenses are adequate to protect eyes from liquids and gases in the lab. a. I and II b. II and IV c. I and III d. III and IV e. II and III

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Chap 02_10e 20. _______________ provides a set of principles intended to help sustain a habitable planet. a. Environmental chemistry b. Analytical chemistry c. Biological chemistry d. Atmospheric chemistry e. Green chemistry 21. Acid washing glassware can replace low concentrations of cations on the surface with H+. Which acids at a concentration of 3–6 M are typically used to clean glassware?

I. acetic acid II. nitric acid III. hydrochloric acid IV. phosphoric acid a. I or II b. II or III c. III or IV d. II or IV e. I or III 22. Which step(s) is/are NOT required in the practice of weighing by difference?

I. Weigh a capped bottle containing the dry reagent.

II. Quickly pour some of the dry reagent from the weighing bottle into a receiver.

III. Weigh the receiver containing the dry reagent.

IV. Recap the bottle that now contains less of the dry reagent.

V. Reweigh the capped bottle that contains less of the dry reagent. a. III b. IV c. V d. IV and V e. I, IV, and V

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Chap 02_10e 23. Which statement(s) is/are TRUE regarding collecting and storing samples for trace analysis?

I. Trace ionic analytes stored in glass are lost by adsorption or contaminated by metals leaching from the glass surface.

II. Plastic bottles are recommended to collect and store ionic analyte samples.

III. Amber glass bottles are best for collecting and storing aqueous samples of organic materials. a. I and II b. II and III c. I and III d. I, II, and III e. I 24. The mass of the empty receiving vessel used with an analytical balance is the a. linearity. b. buoyancy. c. readability. d. tare. e. tolerance.

25. A 0.103 4 M NaOH solution was prepared in the lab at 25°C. By what percentage will the concentration increase if the solution is used in the field at 10°C? The density of water at 10°C is 0.999 702 6 g/mL and the density at 25°C is 0.997 047 9 g/mL.

26. Potassium hydrogen phthalate (KHP) is a primary standard used to determine the concentration of base solutions. The mass of a sample of KHP measured in air is 4.860 7 g. Determine the true mass of KHP (density = 1.636 g/mL). The density of air is 0.001 2 g/mL at 1 bar and 25ºC, and the density of the calibration weights is 8.0 g/ mL.

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Chap 02_10e 27. What mass of CaCl2 must you measure in air to obtain a true mass of 2.811 2 g? The density of air is 0.001 2 g/mL at 1 bar and 25ºC and the density of the calibration weights is 8.0 g/ mL.

28. A solution has a concentration of 0.528 4 M at 13°C. What will its concentration be at 25°C? The density of water at 13°C is 0.999 379 g/mL and the density of water at 25°C is 0.997 047 g/mL.

29. An empty 10-mL volumetric flask weighs 10.271 g. After filling to the mark with distilled water at 20°C, the mass is 20.217 g. What is the true volume of the volumetric flask at 20°C? The density of water at 20°C is 0.998 207 1 g/mL.

30. An empty 10.00-mL volumetric flask has a mass of 11.175 3 g. When filled to the mark with deionized water, the volumetric flask has a mass of 21.178 0 g. Both weighings were performed at 25°C. What is the deviation of the apparent volume from the true volume of the volumetric flask at 25°C? The density of water at 25°C is 0.997 047 9 g/mL.

31. Find the true mass of NaCl (density = 2.16 g/mL) if the apparent mass weighed in air is 25.00 g. The density of air is 0.001 2 g/mL at 1 bar and 25ºC, and the density of the calibration weights is 8.0 g/mL.

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Chap 02_10e 32. A total of 39.56 mL of 0.102 8 M HCl were delivered with a 50 ± 0.05 mL class A buret to neutralize an NaOH solution with an unknown concentration. What is the relative uncertainty associated with the volume delivered by the buret?

33. A weigh bottle has a mass of 10.272 g. When distilled water from a 25-mL transfer pipet is added to the weigh bottle the mass is 35.162 at 25°C. What is the true volume of the 25-mL pipet at 25°C? The density of water at 25°C is 0.997 947 9 g/mL.

34. Describe how to prepare a 20.00-µg/mL iron solution from a 1 000-mg/mL iron standard solution using 10-mL and 50-mL volumetric pipets and 500-mL and 1 000-mL volumetric flasks.

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Chap 02_10e Answer Key 1. c 2. b 3. e 4. a 5. e 6. e 7. b 8. e 9. e 10. c 11. b 12. d 13. a 14. b 15. e 16. c 17. b 18. d 19. c 20. e 21. b 22. a 23. d 24. d 25. 0.266 3%; Calculate the concentration of the solution at 10°C, accounting for thermal expansion. Then find the difference between the concentrations at both temperatures, divide the difference by the concentration at 25°C, and multiply by 100 to determine the percent increase in the concentration.

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Chap 02_10e 26. 4.863 5 g; Use the buoyancy equation to calculate the true mass (m) from the apparent mass (m').

27. 2.810 1 g; Use the buoyancy equation to calculate the apparent mass (m') from the true mass (m).

28. 0.527 2 M; Divide the concentration of the solution at 13°C by the density of water at the same temperature and multiply by the density at 25°C. 29. 9.975 mL; Subtract the mass of the empty flask from the mass of the filled flask to determine the mass of the water in the flask. Divide the mass of water by the density to determine the volume of the water and the flask.

30. 0.032 3 mL; Subtract the mass of the empty flask from the mass of the flask and the water to determine the mass of the water. Divide the mass of the water by its density to determine the volume of the water and the flask. The difference between this volume and 10.00 mL is the deviation in the volume.

31. 25.01 g; Use the buoyancy equation to calculate the true mass (m) from the apparent mass (m').

32. 0.13%; Divide the uncertainty of the buret by the delivered volume and multiply by 100 to determine the relative uncertainty.

33. 24.990 mL; Subtract the mass of the empty bottle from the mass of the filled bottle to determine the mass of the water in the bottle delivered by the pipet. Divide the mass of water by the density to determine the volume of the water and the pipet.

34. Dilute 10 mL of the 1 000-mg/mL iron solution to 500 mL to give a 20 000-µg/mL iron solution. Then, dilute 10 mL of the 20 000-µg/mL iron solution to 1 000 mL to give a 200-µg/mL iron solution. Dilute 50 mL of the 200-µg/mL solution to 500 mL to give a 20-µg/mL solution.

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Chap 03_10e Indicate the answer choice that best completes the statement or answers the question.

1. Blunders may be due to accidental but significant departures form procedure, including a. calculation errors. b. overshooting a titration point. c. dropping, discarding, or contaminating a sample. d. instrument failure. e. all of above.

2. Two students are tasked with determining the milligrams of chloride in a simulated soil sample. Both students extract chloride into aqueous solution and perform triplicate titration with silver nitrate solution and dichlorofluorescein indicator. The students report the following results to their laboratory instructor.

Trial

Student #1 (mg/g)

Student #2 (mg/g)

123

35.9830.1132.88

35.9936.4036.29

Average

32.88

36.23

If the accepted value is 36.90 mg chloride/g soil sample, the laboratory instructor infers that student #1’s work exhibits ___________ student #2’s work. a. higher accuracy and higher precision than b. higher accuracy and lower precision than c. lower accuracy and lower precision than d. lower accuracy and higher precision than e. results impossible to differentiate from 3. Calculate the pH and its uncertainty for a solution that has an H+ concentration of 6.7 (±0.1) × 10−4 M. pH = –log(H+). a. 3.174 (±0.006) b. 3.17 (±0.01) c. 3.2 (±0.1) d. 3.174 (±0.006 5) e. 3 (±1) 4. Calculate the volume of an aquarium that is 73.1 cm long, 30.1 cm deep, and 25.44 cm tall. Report the volume with the correct number of significant figures. a. 55 975 cm3 b. 5.6 × 104 cm3 c. 5.598 × 104 cm3 d. 5.60 × 104 cm3 e. 55 975.9 cm3 Copyright Macmillan Learning. Powered by Cognero.

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Chap 03_10e 5. A Co2+ solution of unknown concentration was titrated with 25.00 mL 0.038 45 M EDTA; the excess EDTA required 42.19 mL 0.020 20 M Zn2+ to reach the end point. How many moles of Co2+ are in the solution?

Co2+ + EDTA → CoEDTA2– Zn2+ + EDTA → ZnEDTA2– a. 1.090 12 × 10−4 mol b. 1.090 1 × 10−4 mol c. 1.090 × 10−4 mol d. 1.09 × 10−4 mol e. 1.1 × 10−4 mol 6. ___________ expresses the margin of uncertainty associated with a measurement. a. Nominal uncertainty b. Maximum uncertainty c. Relative uncertainty d. Minimum uncertainty e. Absolute uncertainty 7. Which statement(s) about significant figures in calculations is/are FALSE? I. Addition and subtraction―round the answer according to the number of decimal places in the number with the most decimal places.

II. Multiplication and division―the answer is limited to the number of digits in the number with the fewest significant digits.

III. Logarithms and antilogarithms―the number of significant figures in the antilogarithm should equal the number of digits in the mantissa. a. I b. II c. III d. I and II e. I and III

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Chap 03_10e 8. For a fixed absolute uncertainty, as the magnitude of the measurement ____________, the percent relative uncertainty ___________. a. increases; decreases b. increases; increases c. decreases; decreases d. decreases; remains constant e. increases; remains constant 9. Which of the paired statements are TRUE for the detection of the systematic error in an analytical method?

I

Analyze a known sample.

II III

Analyze a blank sample.

Analyze the sample using a different analytical technique.

The sample is analyzed by different people in different laboratories with the same or different techniques.

IV

There is no systematic error if the known result is reproduced.

A nonzero answer indicates systematic error.

If the results do not agree, there is systematic error in one or both methods.

Disagreement beyond estimated random error is systematic error.

a. I and II b. II and IV c. I and IV d. II, III, and IV e. I, II, III, and IV 10. Which of the statements are TRUE for the calculation of the uncertainty associated with a molecular mass?

I. The uncertainty in the mass of n identical atoms = n (uncertainty in atomic mass for one atom).

II. Uncertainties for different elements are dependent.

III. The uncertainty for the sum of the masses of different elements is calculated using the equation .

IV. When summing the masses of different elements, the uncertainties for the mass of different elements are propagated as random errors. a. I, II, and III b. I and III c. II, III, and IV d. I and IV e. II and III

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Chap 03_10e 11. Calculate the mass of a concrete slab that is 51.0 cm long, 17.34 cm deep, and 6.2 cm tall with a density of 5.3 g/cm3. Report the mass with the correct number of significant figures. a. 1.0 × 103 g b. 2.9 × 104 g c. 2.91 × 104 d. 1.03 × 103 g e. 2.906 × 104 g 12. ___________ cannot be eliminated, but it may be reduced by better technique. a. Gross error b. Internal error c. Systematic error d. Random error e. Determinate error 13. Calculate the pH of a solution with an H+ concentration of 2.46 10−4 M. Report the pH with the correct number of significant figures. pH = –log[H+] a. 3.6 b. 3.61 c. 3.609 d. 3.609 1 e. 3.609 06 14. ___________ is a consistent error that can be detected and corrected. a. Gross error b. Internal error c. Systematic error d. Random error e. Indeterminate error 15. The distance between two cities is measured. Which reported value of the distance has an ambiguous significant digit? a. 740. km b. 7.40 × 102 km c. 740.0 km d. 740 km e. 7.4 × 102 km

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Chap 03_10e 16. A sample of calcium carbonate is weighed out using the technique weighing by difference. A mass of approximately 10 g is weighed. If a vial containing calcium carbonate powder has an initial mass of 87.36 ± 0.03 g and a final mass of 76.99 ± 0.03 g, what mass of calcium carbonate was taken? a. 10.37 ± 0.04 g b. 10.37 ± 0.03 g c. 10.370 ± 0.042 g d. 10.370 ± 0.004 g e. 10.370 ± 0.005 g 17. Calculate the error for the molecular mass of acetic acid, CH3CO2H. The atomic masses for each element are C = 12.010 6 ± 0.001 0, H = 1.007 98 ± 0.000 14, and O = 15.999 0 ± 0.000 8. a. ± 0.010 u b. ± 0.0026 u c. ± 0.0013 u d. ± 0.0042 u e. ± 0.021 u 18. The tolerance of a Class A 25-mL transfer pipet is ±0.03 mL. If a student uses an uncalibrated Class A 25-mL pipet to deliver a total of 75 mL solution, what is the uncertainty in the delivered 75 mL? a. ±0.09 mL b. ±0.05 mL c. ±0.003 mL d. ±0.03 mL e. ±0.10 mL 19. Calculate the mass of a heterogeneous mixture containing 139.32 g sand, 34.99 g gravel, and 9.372 g salt.

Report the mass with the correct number of significant figures. a. 184 g b. 183.7 g c. 183.68 g d. 184.0 g e. 183.682 g 20. For which measurement is the number of significant digits incorrectly determined? a. 12.983 0 g; six significant digits b. 1 920.3 m; five significant digits c. 0.004 3 g; four significant digits d. 1.003 L; four significant digits e. 59.0 km; three significant digits

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Chap 03_10e 21. Calculate the density of a phosphoric acid solution that is 48.496 g in mass and 28.0 mL in volume. Report the density with the correct number of significant figures. a. 1.7 g/mL b. 1.73 g/mL c. 1.732 g/mL d. 0.577 g/mL e. 0.577 4 g/mL 22. Determine the concentration of Ca2+ in a solution with a pCa2+ of 10.20. Report the concentration with the correct number of significant figures. pCa2+ = –log[Ca2+] a. 6.3 × 10−11 b. 6.31 × 10−11 c. 6.310 × 10−11 d. 6.309 6 × 10−11 e. 6.309 57 × 10−11 23. The tolerance of a Class A 10-mL transfer pipet is ±0.02 mL. A student uses an uncalibrated Class A 10mL transfer pipet to deliver a total of 40 mL of solution. What is the uncertainty in the delivered 40 mL?

24. Calculate the molecular mass of diethyl ether, CH3CH2OCH2CH3. The atomic masses for each element are C = 12.010 6 ± 0.001 0, H = 1.007 98 ± 0.000 14, and O = 15.999 0 ± 0.000 8.

25. The density of a solution is determined using a Class A 25-mL volumetric flask and an analytical balance.

The mass of the empty volumetric flask is 26.987 2 ± 0.000 3 g, the mass of the flask filled with solution is 53.982 0 ± 0.000 3 g, and the tolerance for the uncalibrated 25-mL volumetric flask is ±0.03 mL. What is the density of the solution and its absolute uncertainty?

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Chap 03_10e 26. What is the molarity of a sodium hydroxide solution when 12.90 g of NaOH (FM 40.00) are dissolved in enough water to prepare 0.500 L of solution? Report the molarity to the correct number of significant figures.

27. A 0.199 4-g sample of formic acid (HCOOH) required 42.35 mL of 0.102 3 M NaOH to reach the equivalence point of the titration. What is the formula mass of formic acid? Report the formula mass with the correct number of significant figures.

28. Calculate the formula mass of caffeine (C8H10N4O2), given that C = 12.011 u, H = 1.008 u, N = 14.007 u, and O = 15.999 u.

29. An irregularly shaped solid with a mass of 25.986 3 ± 0.000 3 g displaces 20.5 ± 0.2 mL of water.

Calculate the density of the solid and its absolute uncertainty.

30. A solution has a pNa+ of 9.74 ± 0.12. Calculate [Na+] and its uncertainty. pNa = −log[Na+]

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Chap 03_10e 31. A 25.00-mL transfer pipet consistently delivers 0.12% less than its nominal value. What is the systematic error of the pipet?

32. Calculate the number of moles of NaCl in 10.00 ± 0.04 mL of a 1.05 ± 0.02 M NaCl solution. Report the absolute and percent relative uncertainty.

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Chap 03_10e Answer Key 1. e 2. c 3. a 4. d 5. c 6. e 7. a 8. a 9. e 10. d 11. b 12. d 13. c 14. c 15. d 16. a 17. b 18. a 19. c 20. c 21. b 22. a 23. ±0.08 mL 24. 74.121 ± 0.0043 u 25. 1.080 ± 0.001 g/mL 26. 0.654 M Copyright Macmillan Learning. Powered by Cognero.

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Chap 03_10e 27. 46.03 g/mol 28. 194.194 u 29. 1.27 ± 0.01 g/mL 30. 1.8 (±0.5) 10−10 M 31. 0.03 mL 32. 10.5 ± 0.2 mmol; 1.9%

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Chap 04_10e Indicate the answer choice that best completes the statement or answers the question.

1. To determine if two methods for measuring the amount of calcium in milk produce results that are statistically different, a technician analyzes a milk sample five times with each method. If the F test determines the standard deviations of the two methods are significantly different, which set of equations must the technician use to calculate t? a.

, where

b.

, where

c.

, where degrees of freedom =

d.

, where degrees of freedom =

e. None of these sets of equations can be used to calculate t.

2. A pair of students standardize a sodium hydroxide solution using KHP and obtain the following concentrations: 0.102 8 M, 0.103 1 M, 0.120 0 M, 0.103 0 M, and 0.102 6 M. The students suspect that 0.120 0 M is an outlier. Is 0.120 0 M an outlier and what is the average concentration of the sodium hydroxide solution? Refer to Table 4-6 for critical values of G. a. 0.120 0 M is an outlier; = 0.102 9 M b. 0.120 0 M is not an outlier;

= 0.106 3 M

c. 0.120 0 M is not an outlier;

= 0.102 9 M

d. 0.120 0 M is an outlier;

= 0.106 3 M

e. It is not possible to determine if 0.120 0 M is an outlier.

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Chap 04_10e 3. The Gaussian distribution is characterized by the ___________ and the ___________. a. mean; precision b. accuracy; standard deviation c. mean; standard deviation d. accuracy; precision e. precision; standard deviation 4. Doubling the number of calibration curve data points decreases the standard uncertainty in the slope and yintercept by a factor of a. 2. b. ½. c. 4. d. 2. e. 1/2 5. The density of a solution is measured six times with results of 1.098, 1.100, 1.089, 1.095, 1.097, and 1.101 g/mL. Calculate the 95% confidence interval for the density. Refer to Table 4-4 for values of Student’s t. a. 1.0967 ± 0.0043 g/mL b. 1.0967 ± 0.0038 g/mL c. 1.0967 ± 0.0041 g/mL d. 1.0967 ± 0.0045 g/mL e. 1.0967 ± 0.0039 g/mL 6. Which statements regarding outliers are TRUE? I. An outlier is a datum that is far from other points.

II. All outliers must be discarded before averaging the rest of the data.

III. Using statistical tools to discard outliers is a poor substitute for good record keeping.

IV. The Q test is a mathematically simpler but more limited test for outliers than is the Grubbs test. a. I and IV b. I and II c. I and III d. I, III, and IV e. I, II, III, and IV

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Chap 04_10e 7. Which statements regarding the method of least squares are true?

I. The method of least squares draws the “best” straight line through experimental data points.

II. The equation for a straight line is y = bx + m, where m is the slope and b is the y-intercept.

III. Deviations in y are minimized because the uncertainty in y is assumed to be greater than the uncertainty in x.

IV. The slope and y-intercept are calculated using determinants. a. I, II, and III b. I, III, and IV c. I, II, and IV d. II, III, and IV e. III and IV 8. What is the null hypothesis for the comparison of two means using a t test? a. The two sets of measurements come from populations with the same standard deviations. b. The two sets of measurements come from populations with different population means. c. The two sets of measurements come from populations with the same population mean and different population standard deviations. d. The two sets of measurements come from populations with different standard deviations. e. The two sets of measurements come from populations with the same population mean.

9. A lab technician was asked to determine the weight percent of Fe in an ore sample. The following results were obtained for the six analyses the technician did. What are the degrees of freedom of this data set?

8.25%, 8.28%, 8.03%, 8.19%, 8.33%, 8.24% a. 1 b. 5 c. 6 d. 7 e. 4

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Chap 04_10e 10. A technician determines the concentration of calcium in milk using two instrumental methods. If Fcalculated > Ftable for the two sets of calcium data, what conclusion(s) can the technician make?

I. The difference in standard deviations for the two instrumental methods is significant.

II. The difference in standard deviations for the two instrumental methods is not significant.

III. The data come from populations with the same standard deviation.

IV. The data do not come from populations with the same standard deviation. a. I and III b. I and IV c. II and III d. II and IV e. II 11. Which term is INCORRECTLY defined? a. linear range – the range over which the response of an analytical method is proportional to analyte concentration b. corrected absorbance – the average absorbance of the blank measurements is subtracted from each measured absorbance c. blank solution – a solution containing all reagents and solvents used in the analysis to which no analyte is deliberately added d. standard solutions – solutions for which the standard deviation is known e. dynamic range – concentration range over which there is a measurable response to analyte, though the response is not linear 12. An ore sample was analyzed for its Fe content. Student A analyzed the sample a total of six times and her results had a standard deviation of 1.33. The same sample was analyzed five times by Student B and his results had a standard deviation of 3.42. To determine if their standard deviations are similar, they perform an F test. The calculated F value is ________, and the table value of F is _________. Refer to Table 4-3 for critical values of F. a. 6.61; 7.39 b. 2.57; 7.39 c. 6.61; 9.36 d. 0.151; 9.36 e. 6.61; 6.98

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Chap 04_10e 13. A data point may be excluded when I. Gcalculated is greater than Gtable.

II. Gcalculated is less than Gtable.

III. your TA or supervisor says to disregard the datum.

IV. the datum is the result of a faulty procedure. a. II and IV b. I, III, and IV c. I and IV d. II, III, and IV e. I, II, III, and IV 14. Rose gold is a gold alloy that contains a small amount of copper. The copper content of rose gold was measured by two methods for six different samples collected from a jewelry production plant.

Sample

Method 1

Method 2

A

2.2%

2.5%

B

1.8%

2.2%

C

3.0%

3.2%

D

2.5%

3.0%

E

2.6%

2.3%

F

2.6%

2.7%

Which test would you use to determine if the results obtained by the two methods agree within experimental error? a. Grubbs test b. F test c. paired t test to compare individual difference d. t test to compare mean values of two sets of replicate measurements e. compare a measured value to a “known” value

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Chap 04_10e 15. A biosensor was developed to measure the concentration of aqueous arsenic. To determine the accuracy of the sensor, the sensor was tested against a standard reference solution of known concentration. The concentration measured by the new biosensor can be considered accurate if a. the accepted, known concentration of the reference solution falls outside the 95% confidence interval for the biosensor results. b. the accepted, known concentration of the reference solution falls within the 95% confidence interval for biosensor results. c. the standard deviations for the concentration determined by the biosensor and the accepted, known concentration of the reference solution are not significantly different. d. the standard deviations for the concentration determined by the biosensor and the accepted, known concentration of the reference solution are not significantly different, and t calculated < t table. e. the standard deviations for the concentration determined by the biosensor and the accepted, known concentration of the reference solution are not significantly different, and t calculated < t table.

16. Which statements regarding the mean and standard deviation are TRUE? I. As the number of measurements increases,

approaches μ if there is no random error.

II. The square of the standard deviation is the average deviation.

III. The mean is the center of the Gaussian distribution.

IV. The standard deviation measures the width of the Gaussian distribution. a. II, III, and IV b. I, III, and IV c. I and II d. I, II, and IV e. III and IV 17. Which statement(s) regarding confidence intervals is/are TRUE? I. As the percentage confidence increases, the confidence interval range decreases.

II. Confidence intervals are calculated using the calculated mean and standard deviation of a set of n measurements and the result of the F test.

III. The 95% confidence interval would include the true population mean in 95% of the sets of n measurements. a. I b. II c. I and III d. I and II e. III

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Chap 04_10e 18. Student’s t is a statistical tool used most frequently to I. express confidence intervals.

II. compare results from different experiments.

III. evaluate the probability of an experimental value agreeing with a “known” value.

IV. determine if a questionable data point should be discarded. a. I b. II c. I and II d. I, II, and III e. I, II, III, and IV 19. A technician determines the concentration of calcium in milk using two instrumental methods. If Fcalculated < Ftable for the two sets of calcium data, what conclusion(s) can the technician make?

I. The difference in standard deviations for the two instrumental methods is significant.

II. The difference in standard deviations for the two instrumental methods is not significant.

III. The data come from populations with the same standard deviation.

IV. The data do not come from populations with the same standard deviation. a. I and III b. I and IV c. II and III d. II and IV e. I 20. Calculate the mean and standard deviation for the following results for the concentration of lead in a soil sample.

23.2 ppm, 20.1 ppm, 24.7 ppm, 19.9 ppm, 21.8 ppm a. 21.9 3.4 ± (n = 5) ppm b. 21.9 2.0 ± (n =5) ppm c. 27.4 2.0 ± (n = 5) ppm d. 22.0 1.8 ± (n = 5) ppm e. 27.4 4.2 ± (n = 5) ppm

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Chap 04_10e 21. The absorbance for a dye sample of unknown concentration is measured and the corrected absorbance is 0.5482. Calculate the concentration of the dye sample using the calibration curve that was prepared with a series of standard dye solutions of known concentration.

a. 0.000 65 μg/mL b. 0.000 52 μg/mL c. 0.000 92 μg/mL d. 0.0011 μg/mL e. 0.000 79 μg/mL 22. Use the Grubbs test to determine whether 40.73% should be discarded from the data set 40.73, 38.43, 37.23, 38.66, 39.04, 37.66, and 38.32% obtained in the determination of CaO in a chalk sample. Refer to Table 4-6 for critical values of G.

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Chap 04_10e 23. A student used two different methods to measure the sulfite content (wt%) in several solid samples. The results are shown in the table.

Sample Method A Method B 1 48.17 47.68 2 43.26 43.86 3 49.38 50.12 4 51.82 50.85 5 44.95 44.01 6 51.27 50.38 Determine if the two methods give significantly different results at the 95% confidence level. Refer to Table 4-4 for values of Student’s t.

24. The molarity of a sodium hydroxide solution was determined by titration against KHP to be 0.102 5 M, 0.108 7 M, 0.110 0 M, 0.095 1 M, and 0.099 7 M. Calculate the 95% confidence level for the molar concentration of the sodium hydroxide solution. Refer to Table 4-4 for values of Student’s t.

25. The molarity of a sodium hydroxide solution was determined by titration against KHP. The results from one lab section were 0.102 5 M, 0.108 7 M, 0.110 0 M, 0.105 2 M, and 0.099 7 M. The lab instructor suspects that 0.099 7 M is an outlier. Can the lab instructor throw out the data point? Refer to Table 4-6 for critical values of G.

26. The average percent purity for a batch of acetic acid manufactured by Acetic Acid Corp is 97.5% with a standard deviation of 0.2%. The minimum percent purity a buyer requires is 97.0%. What percentage of batches will have a percent purity unacceptable to the customer? Refer to Table 4-1 for the area under each portion of a Gaussian curve.

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Chap 04_10e 27. A technician determines the concentration of calcium in milk using two instrumental methods. For method A, ppm and s = 0.40 (n = 8) and for method B, ppm and s = 0.72 (n = 6). Is the difference in standard deviations significant at the 95% confidence level? Refer to Table 4-3 for critical values of F.

28. A smartphone-assisted colorimetric study of permanganate standard solutions in the concentration range of 0.08–0.40 mM produced a calibration curve for the G-values measured. The quadratic equation fitted to the data is y = 497.27x 2 – 791.33x + 259.44. If an unknown had a G-value of 108, what was the concentration of permanganate in the unknown?

29. A researcher has developed a new analytical method to determine the percent by mass iron in solids. To test the new method, the researcher purchases a standard reference material sample that is 2.85% iron by mass. Analysis of the iron standard with the new method returns values of 2.75%, 2.89%, 2.77%, 2.81%, and 2.87%. Does the new method produce a result that is significantly different from the standard value at the 95% confidence level? Refer to Table 4-4 for values of Student’s t.

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Chap 04_10e 30. The concentration and absorbance data in the table were collected for compound X. The least squares line for the plot of absorbance versus concentration is y = 6.83x + 0.088 1. A sample of compound X of unknown concentration has a corrected average absorbance of 0.652 1 (three replicate measurements).

Calculate the concentration of compound X in the sample and the uncertainty in the concentration.

Concentration (mg/mL) Absorbance 0.050 0.442 6 0.075 0.574 4 0.100 0.784 1

31. Two lab sections were required to determine the amount of zinc in a multivitamin. The first lab section with 16 students reported an average value with a standard deviation of 1.86. The second lab section with 13 students reported an average value with a standard deviation of 1.43. Is the difference in standard deviations significant at the 95% confidence level? Refer to Table 4-3 for critical values of F.

32. The content of Mn in a solid sample was determined using smartphone calorimetry and spectrophotometry.

Smartphone colorimetry gave a result of 0.70 ± 0.04% (n = 6) and spectrophotometry gave a result of 0.69 ± 0.01% (n = 6). Determine if the results of the two methods differ significantly at the 95% confidence level. Refer to Table 4-4 for values of Student’s t.

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Chap 04_10e Answer Key 1. c 2. a 3. c 4. e 5. d 6. d 7. b 8. e 9. b 10. b 11. d 12. a 13. c 14. e 15. b 16. e 17. e 18. d 19. c 20. b 21. a 22. Gcalculated = 1.909 < Gtable = 1.938, so the data point should not be discarded.

23. t calculated = 1.006 < t table = 2.571, so the results of the two methods are not significantly different at the 95% confidence level. Use a paired t test to compare the differences between the two methods for each individual measurement.

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Chap 04_10e 24. 0.1032 ± 0.0077 M; Calculate the average (0.103 2 M) and standard deviation (0.006 22 M) of the five values. At the 95% confidence level with 4 degrees of freedom, t = 2.776. The confidence interval is calculated using the equation .

25. For the lab section, , s = 0.004 3 M, n = 5. The data point cannot be thrown out because Gcalculated(= 1.29) is smaller than Gtable(= 1.67).

26. 0.62%; The value of z is −2.5 with an area of 0.493 8. The percentage below 97% is 0.500 0 − 0.493 8 = 0.006 2 or 0.62%.

27. Fcalculated = (0.72)2/(0.40)2 = 3.24; Ftable = 3.97 for n1 − 1 = 5 and n2 − 1 = 7 degrees of freedom; Fcalculated(= 3.24) < Ftable(= 3.97), so the standard deviations are not significantly different.

28. 0.222 mM; Insert the value of 108 into the quadratic equation for y and solve for x.

29. For the new method, , s = 0.078%, and t = 2.776 for 4 degrees freedom at 95% confidence and n = 5. At the 95% confidence level, the confidence interval for the new method is 2.798 ± 0.097% by mass Fe, or 2.70% to 2.90%. The accepted value falls within the 95% confidence interval, so the new method produces a result that is not significantly different.

30. 0.0826 ± 0.0039 mg/mL; Use the equation of the least squares line to determine the concentration from the absorbance. Then use the provided equations to calculate the uncertainty.

31. Fcalculated = 1.69 < Ftable = 3.18; the standard deviations are not significantly different.

32. t calculated = 0.58 < t table = 2.228, so the results of the two methods are not significantly different at the 95% confidence level. Use the F test to determine if the standard deviations are significantly different. Then use the appropriate set of equations to compare the results of the two different sets of data.

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Chap 05_10e Indicate the answer choice that best completes the statement or answers the question.

1. A _________ is an analytical result that indicates the sample concentration is below a stated limit when the concentration is actually above the limit. a. positive result b. false negative c. negative result d. false positive e. negative 2. Which statements regarding internal standards are TRUE? I. The response factor is calculated by using a mixture containing known amounts of analyte and internal standard and the detector response to the analyte and the internal standard.

II. Internal standards are useful where instrument response or quantity of sample analyzed varies over time.

III. Detector response is inversely proportional to concentration.

IV. Detector response to analyte and internal standard is assumed to be constant over a range of concentrations. a. I, II, and III b. II, III, and IV c. I and II d. I and III e. I, II, and IV 3. A(n) ________________ is a known amount of compound, different from the analyte, that is added to the unknown. a. standard addition b. internal standard c. external standard d. self-contained standard e. known standard 4. A solution is 2.00% by mass of compound A and 3.50% by mass of internal standard S. Separation by gas chromatography returns peak areas of 38 234 for compound A and 32 880 for internal standard S. Calculate the response factor. a. 0.664 b. 2.03 c. 1.50 d. 0.491 e. 1.17

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Chap 05_10e 5. Which of the studies is NOT required for method validation? a. method specificity b. linearity c. range d. calibration e. limit of detection 6. ____________ is the extent to which a method can distinguish the analyte from everything else in the sample. a. Sensitivity b. Selectivity c. Limit of detection d. Limit of quantitation e. Linearity 7. Specifications might include a. sampling requirements. b. accuracy and precision. c. selectivity and sensitivity. d. recovery of fortification. e. all of above.

8. A _________ is a sample that has been exposed to the site of sampling, containing all components except analyte, and it is taken through all steps of the analytical procedure. a. method blank b. reagent blank c. field blank d. spike blank e. matrix blank

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Chap 05_10e 9. The response factor, F, for the analysis of pentane using the internal standard decane is 3.5. Which statement(s) about the analysis is/are TRUE? I. The pentane concentration is 3.5 times greater than the decane concentration.

II. The decane concentration is 3.5 times greater than the pentane concentration.

III. Detector response to pentane is 3.5 times greater than detector response to decane.

IV. Detector response to decane is 3.5 times greater than detector response to pentane. a. I and III b. III c. IV d. II and IV e. II 10. A set of five solutions are prepared by delivering 10 mL of a solution with an unknown concentration and increasing volumes of standard (0, 5, 10, 15, and 20 mL) into 50-mL volumetric flasks. Analytical reagents are then added and diluted to volume. The x-intercept of the best-fit line is −3.90 mM. What is the concentration of the unknown in the original solution? a. 39.0 mM b. 0.780 mM c. 19.5 mM d. 3.90 mM e. 195 mM 11. The ______________ is everything in the sample, except analyte. a. blind sample b. performance test sample c. spike d. matrix e. assessment 12. A _________ is a sample containing all components except analyte, and it is taken through all steps of the analytical procedure. a. method blank b. reagent blank c. field blank d. spike blank e. matrix blank

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Chap 05_10e 13. ____________ is the concentration range over which there is measurable instrument response. a. Accuracy range b. Dynamic range c. Linear range d. Detection range e. Range 14. Which statement regarding standard addition to a single solution is FALSE? a. Analyte signal is measured after each standard addition. b. The volume of standard added must be kept small to minimize altering the matrix. c. Detector response is corrected to take into account dilution before being plotted on the y-axis. d. The absolute value of the y-intercept for the best-fit line is the unknown analyte concentration. e. [S]i(Vs /Vi) is plotted on the x-axis.

15. There are several types of precision involved with method validation. Which of the precisions is NOT associated with method validation? a. instrument precision b. final precision c. levelofdifficulty: Intermediate precision d. interlaboratory precision e. intra-assay precision 16. ______________ are concentrations or amounts found by applying a calibration procedure to the _________________. a. Raw data; treated data b. Results; treated data c. Treated data; raw data d. Treated data; results e. Results; raw data

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Chap 05_10e 17. Which of the following is/are NOT a way to express the precision of a method?

I. standard deviation II. standard uncertainty III. confidence interval IV. nearness to the true value a. IV b. I and III c. II and III d. III e. III and IV 18. Which of the following is NOT part of the process of assessing the linearity of a calibration? a. visually inspecting the scatter of the data about the least-squares line b. preparing a residual plot c. determining the square of the correlation coefficient, R2 d. assessing how close the y-intercept is to 0 e. All of the steps are used to assess the linearity of a calibration.

19. Which of the following is/are NOT a way to demonstrate the accuracy of a method?

I. Determine the limit of detection for the method II. Analyze a certified reference material in a matrix similar to the sample III. Compare the results of two or more different analytical methods IV. Analyze a blank sample spiked with a known amount of analyte in the same matrix as the sample a. IV b. II and III c. I and III d. III e. I 20. Which statement regarding the multiple solutions with constant volume standard addition method is FALSE? a. Equal volumes of unknown are pipetted into a set of volumetric flasks with identical volume. b. Equal volumes of standard are pipetted into each volumetric flask. c. Analytical reagents are added to each volumetric flask and diluted to volume. d. Detector response is plotted on the y-axis versus the concentration of the diluted standard on the x-axis. e. The absolute value of the x-intercept for the best-fit line is the concentration of the diluted unknown.

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Chap 05_10e 21. The ______________ state(s) what steps will be taken and how they will be carried out. a. assessment b. calibration check c. control charts d. specifications e. standard operating procedure 22. Which practice is NOT used to gauge accuracy? a. raw data and results from calibration checks b. fortification recoveries c. quality control samples d. blank samples e. sample replications 23. The percent by mass octane is quantified for an organic sample via gas chromatography using decane as the internal standard. A vial containing 0.502 3 g organic sample and 0.238 4 g decane is prepared and analyzed, returning peak areas of 119 993 for octane and 83 755 for decane. If a response factor of 2.50 from previous work analyzing octane with a decane internal standard is used, what is the percent by mass octane in the organic sample?

24. The solvent 2-propanol contained a trace amount of water, and the amount of water was determined using gas chromatography with a thermal conductivity detector (GC-TCD) utilizing methanol as the internal standard. A 4.900-mL sample of the solvent was mixed with 0.100 mL methanol. After the GCTCD analysis of the mixture, the peak areas of water and methanol were found to be 15.79 and 27.77, respectively. A 25 mL aliquot of water was then added to the above solution. The new peak areas for water and methanol were measured to be 18.96 and 21.43, respectively. Find the percent water by volume in the 2-propanol solvent, assuming all volumes were additive.

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Chap 05_10e 25. A chemist is tasked with determining the concentration of nickel in a water sample drawn from a well next to a nickel mine. To analyze the sample using atomic absorbance spectroscopy (AAS), she prepares six samples (0–5) by adding 10.00 mL of well water to six 50.00-mL volumetric flasks. She then adds increasing volumes of 25.00-ppm nickel solution to the volumetric flasks, and dilutes the resulting solutions to 50.00 mL with 0.1 M HNO3. Following AAS analysis, a graph of absorbance versus the nickel standard concentration in parts per million (ppm) gave a best-fit line of y = 0.0428x + 0.0487.

Calculate the ppm nickel concentration in the well water sample.

26. Amperometry can be applied to the analysis of dopamine, a hormone and neurotransmitter. Signals from seven replicate samples with a concentration close to three times the detection limit were 1.26, 1.18, 1.13, 1.12, 1.22, 1.19, and 1.11 nA. The average signal for ten reagent blanks was 0.17 nA. What is the signal detection limit?

27. A 3.42-ppb Ni2+ concentration was found in an electroplating waste sample. To a 9.00-mL aliquot of the waste sample, 1.00 mL of a 1.25-ppb Ni2+ standard was added. Analysis of the spiked sample gave a concentration of 3.20-ppb Ni2+. Find the percent recovery of the spike.

28. The content of ethanol in alcoholic beverages can be quantified by gas chromatography using acetonitrile as an internal standard. First, 1.000 mL of a 20.0% ethanol standard was mixed with 0.100 mL acetonitrile, and then diluted to 5.000 mL with deionized water. The peak area of ethanol was measured to be 4 528, and that of acetonitrile was determined to be 3 973. Later, 1.000 mL of an alcoholic beverage of unknown concentration was mixed with 0.100 mL acetonitrile and diluted to 5.000 mL with deionized water. Peak areas were 3 395 for ethanol, and 4 109 for acetonitrile. What is the concentration of ethanol in the alcoholic beverage?

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Chap 05_10e 29. Epigalocatechin gallate (EGCG) is the primary catechin found in green tea extract. The high-performance liquid chromatography–ultraviolet (HPLC-UV) analysis of green extract gave a signal of 72 for EGCG. Later, 5.00 mL of green tea extract and 1.00 mL of a 200 mM standard EGCG were mixed together and diluted to 10.00 mL. The analysis of the spiked sample gave a signal of 89. Calculate the concentration of EGCG in the green tea extract.

30. A water sample collected downstream from a gold-mining operation returns a cyanide concentration of 6.18 μg/L CN−. A 10-mL aliquot of the water sample is spiked with 100 μL of 500.0 μg/L sodium cyanide and analyzed. Analysis of the spiked sample returns a cyanide concentration of 10.5 μg/L CN−.

What is the percent recovery for the spike?

31. Eight absorbance readings from replicate caffeine samples with a concentration approximately three times the estimated detection limit are 0.123 9, 0.109 8, 0.118 2, 0.115 0, 0.121 9, 0.125 1, 0.119 9, and 0.109 9. The average absorbance of eight reagent blanks is 0.018 2. The slope of the calibration curve for higher concentrations is 0.590 mM−1. What is the detection limit?

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Chap 05_10e Answer Key 1. b 2. e 3. b 4. b 5. d 6. b 7. e 8. c 9. d 10. c 11. d 12. a 13. b 14. d 15. b 16. c 17. a 18. e 19. e 20. b 21. e 22. e 23. 27.2% octane 24. 0.93% water 25. 5.69 ppm 26. 0.34 nA Copyright Macmillan Learning. Powered by Cognero.

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Chap 05_10e 27. 97.6% 28. 14.5% 29. 27.2 mM 30. 88.5% recovery when accounting for dilution of the original sample 31. 0.030 mM

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Chap 06_10e Indicate the answer choice that best completes the statement or answers the question.

1. For a reaction with a ΔHo = +50 kJ and a ΔSo = −100.0 J/K, ΔGo is ____________ and the equilibrium constant is ____________. a. negative; less than 1 b. negative; greater than 1 c. positive; less than 1 d. positive; greater than 1 e. zero; 1 2. For the reaction HCN(aq) + H2O(l) ⇌ CN−(aq) + H3O+(aq), which action(s) will NOT shift the equilibrium toward the products?

I. Adding HCN II. Adding H2O III. Removing CN− IV. Removing H3O+ a. I b. II c. I, III, and IV d. II and IV e. III 3. Nickel forms three complexes with hydroxide, NiOH+(aq), Ni(OH)2(aq), and Ni(OH (aq). Which statement(s) regarding the impact of complex formation on the solubility of nickel hydroxide are FALSE? I. Hydroxide behaves as a Lewis acid when forming a complex with nickel.

II. The total nickel concentration is [Ni2+]tot = [Ni2+] + [NiOH+] + [Ni(OH)2] + [Ni(OH ].

III. As [OH−] increases, the solubility of Ni(OH)2 initially decreases but then increases.

IV.

a. I b. II and III c. III d. IV e. I, II, III, and IV Copyright Macmillan Learning. Powered by Cognero.

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Chap 06_10e 4. Which statements are TRUE for enthalpy and entropy?

I. Enthalpy is the heat flow, q, under constant applied pressure.

II. Standard conditions for entropy and enthalpy changes are 1 M concentration, 1 bar pressure, and 298.15 K. III. The entropy change for a system is the amount of energy, at a given temperature, that is dispersed into motion of the molecules in the system.

IV. When the value of ΔH is negative, heat is released. When the value of ΔS is negative, entropy is increased. a. I, II, and III b. I, III, and IV c. II and IV d. III and IV e. Another combination of statements is true.

5. A student attempts to use Pb(NO3)2 to separate Cl− and I− from a solution that is 0.05 M Cl− and 0.10 M I−. As Pb(NO3)2 is added, which anion will precipitate first and what will the concentration of the anion be when the second anion begins to precipitate? Ks p = 1.7 × 10−5 for PbCl2 and Ks p = 9.8 × 10−9 for PbI2. a. I− precipitates first; 0.006 8 M I− b. I− precipitates first; 0.001 2 M I− c. I− precipitates first; 2.88 × 10−5 M I− d. Cl− precipitates first; 0.001 2 M Cl− e. Cl− precipitates first; 0.10 M Cl− 6. Which statement about Brønsted-Lowry acids and bases is INCORRECT? a. Brønsted-Lowry acids are proton donors. b. is the conjugate acid of the weak base NH3. c. When an acid and a base react, the acid and base neutralize each other and form a salt. d. Water reacts with itself to form hydronium and hydroxide, a process called autoprotolysis. e. CH3CN is a protic solvent.

7. In which solution is the solubility of silver chloride the greatest? a. pure water b. 0.05 M CaCl2 c. 0.10 M NaCl d. 0.10 M FeCl3 e. 0.15 M AgNO3 Copyright Macmillan Learning. Powered by Cognero.

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Chap 06_10e 8. Which of the following compounds is NOT an aprotic solvent? a. ethanol, CH3CH2OH b. acetonitrile, CH3CN c. dimethyl ether, CH3OCH3 d. cyclohexane, C6H12 e. ethyl acetate, CH3CO2C2H5 9. For a triprotic acid, which of the following is NOT true? a. Ka1Kb3 = Kw b. c. Ka1Kb2 = Kw d. Ka3Kb1 = Kw e.

10. What is the pH of a solution if the OH− concentration is 5.0 ×10−5 M? a. 5.70 b. 4.30 c. 9.70 d. 5.00 e. 9.00 11. When the direction for a reaction is reversed, the new value of K is a. K´ = log(K). b. K´ = ln(K). c. K´ = 10K. d. K´ = 1/K. e. K´ = eK. 12. Which class of acids or bases and strength have been incorrectly matched? a. RCO2H: weak acid b. R3N: weak base c.

: weak base

d. R3NH+: weak acid e. R4NOH: weak base Copyright Macmillan Learning. Powered by Cognero.

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Chap 06_10e 13. The pH of a solution is 9.65. The solution is ________and the hydroxide concentration is ________ than the hydronium concentration. a. acidic; less b. acidic; greater c. neutral; equal d. basic; greater e. basic; less 14. Which acid is the strongest? a. acetic acid, pKa = 4.756 b. benzoic acid, pKa = 4.202 c. cyanoacetic acid, pKa = 2.472 d. phenol, pKa = 9.997 e. formic acid, pKa = 3.744 15. Determine the value of K for the reaction formed when the two given reactions are added together.

CO2(aq) + H2O(l) ⇌ (aq) + H+(aq) K1 = 4.5 × 10−7 (aq) ⇌ H+(aq) + (aq) K = 4.7 × 10−11 2

_________________________________________________________________________________________________

CO2(aq) + H2O(l) ⇌ 2H+(aq) +

(aq) K3 = ?

a. 4.5 × 10−7 b. 4.7 × 10−11 c. 2.1 × 10−17 d. 9.6 × 103 e. 1.0 × 10−4 16. What is the molar solubility of La(IO3)3 in pure water? Ks p = 1.0 × 10−11 for La(IO3)3. a. 3.2 × 10−6 M b. 1.8 × 10−6 M c. 7.8 × 10−4 M d. 1.1 × 10−6 M e. 1.4 × 10−3 M

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Chap 06_10e 17. Calculate ΔGo for the reaction Ag+ + 2

⇌

, for which K = 4.36 at 25oC.

a. −3 650 J b. 3 650 J c. −306 J d. 306 J e. −1 585 J 18. The molar solubility of a saturated iron(II) carbonate solution derived from the Ksp value is 5.6 × 10−6 M. The molar solubility is greater than 5.6 × 10−6 M when accounting for additional reactions. Which of the reactions will not increase the solubility of iron(II) carbonate? a. Fe2+(aq) + OH−(aq) ⇌ FeOH+(aq) b. CO32−(aq) + H+(aq) ⇌ HCO3−(aq) c. HCO3−(aq) + H+(aq) ⇌ H2CO3(aq) d. H2CO3(aq) ⇌ CO2(g) + H2O(l) e. FeOH+(aq) + OH−(aq) ⇌ Fe(OH)2(s) 19. All of the following are strong acids, EXCEPT a. HCl. b. HNO3. c. HBr. d. HClO4. e. HF. 20. For the reaction aA(aq) + bB(s) ⇌ cC(l) + dD(g), which statement is FALSE? a. b. The concentrations of solutes are expressed in moles per liter. c. The concentrations of gases are expressed in bars. d. The concentrations of pure solids, pure liquids, and solvents are omitted because they are negligible. e. The quantities in the equilibrium expression are the ratio of the concentration of the species to the concentration in its standard state.

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Chap 06_10e 21. Ag+ is being considered as a reagent for separating IO3− from in a solution that is 0.060 M in K2CO3 and 0.070 M in NaIO3. Which anion will precipitate first and what will its concentration be when the second anion begins to precipitate? Can the two ions be separated with 99.99% efficiency? Ks p = 3.1 × 10−8 for AgIO3; Ks p = 8.1 × 10−12 for Ag2CO3

22. Calculate the pH of water at 50°C. Kw = 5.31 × 10−14 at 50°C.

23. For the reaction NH3(aq) + H2O(l) the following equilibrium?

2 (aq) + 2OH−(aq)

(aq) + OH−(aq), K = 1.8 × 10−5. What is the K value for

2NH3(aq) + 2H2O(l)

24. The Ka value for CH3CH2NO2 is 2.7 × 10−9. Calculate Kb for

.

25. Calculate the molar solubility for calcium fluoride in a 0.10 M NaF solution. Ks p = 3.2 × 10−11 for CaF2.

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Chap 06_10e 26. What is the molar solubility of brucite crystal, Mg(OH)2, in pure water. Ks p = 7.1 × 10−12 for brucite crystal, Mg(OH)2.

27. Calculate the total nickel concentration for a solution prepared by dissolving Ni(OH)2(s) in a 0.10 M OH− solution. Ks p = 6 × 10−16, b 1 = 1.26 × 104, b 2 = 1 × 109, and b 3 = 1 × 1012.

28. A reaction has a Δ

= −167.5 kJ/mol and Δ

= 57.3 J/mol K. Calculate the equilibrium constant

for the reaction.

29. Calculate the pH for a solution with an OH− concentration of 0.293 M.

30. Oxalic acid, HO2CCO2H, is a diprotic acid with Ka1 = 5.62 × 10−2 and Ka2 = 5.42 × 10−5. Calculate Kb for .

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Chap 06_10e Answer Key 1. c 2. b 3. a 4. a 5. b 6. e 7. a 8. a 9. c 10. b 11. d 12. e 13. d 14. c 15. c 16. c 17. a 18. e 19. e 20. d 21. Iodate precipitates first. [IO3−] = 2.67 × 10−3 M when carbonate begins to precipitate. The percentage of iodate removed is 96.2%, so they cannot be successfully separated with 99.99% efficiency.

22. 6.637 23. 3.1 × 109 24. 3.7 × 10−6 25. 3.2 × 10−9 M Copyright Macmillan Learning. Powered by Cognero.

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Chap 06_10e 26. 1.2 × 10−4 M 27. 6.06 × 10−5 M; [Ni2+]tot = [Ni2+] + [NiOH+] + [Ni(OH)2] + [Ni(OH] = Ks p/[OH−]2 + b 1[Ni2+][OH−] + b 2[Ni2+][OH−]2 + b 3[Ni2+][OH−]3 Solve for [Ni2+] from Ks p when [OH−] = 0.10 M. Then substitute for b 1, b 2, and b 3.

28. K = 2.18 × 1032 29. 13.467; Calculate pOH and subtract from 14.

30. 1.78 × 10−13

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Chap 07_10e Indicate the answer choice that best completes the statement or answers the question.

1. A mixture of Cl−, Br−, and I− is titrated by AgNO3. The order of precipitation of the ions from first to last is _______________. Ks p = 1.8 × 10−10 for AgCl, Ks p = 5.0 × 10−13 for AgBr, and Ks p = 8.3 × 10−17 for AgI. a. Cl−, Br−, I− b. Br−, I−, Cl− c. I−, Cl−, Br− d. Cl−, I−, Br− e. I−, Br−, Cl− 2. Water hardness can be determined by titrating a water sample with EDTA in the presence of pH 10 NH3– NH4Cl buffer and calmagite as the indicator. A 25.00-mL water sample required 23.18 mL of 0.0102 4 M EDTA to reach the end point. Report the water hardness in mg CaCO3/L. a. 950.0 mg CaCO3/L b. 380.6 mg CaCO3/L c. 1 105 mg CaCO3/L d. 442.6 mg CaCO3/L e. 619.5 mg CaCO3/L 3. The end point of some precipitation titrations for mixtures can be up to 3% high. The error is attributed to a. ion pairs. b. coprecipitation. c. turbidity. d. solubility of the precipitates. e. redox chemistry.

4. Which of the terms is INCORRECTLY defined? a. titration error: the difference between the end point and the equivalence point b. direct titration: titrant is added to analyte until reaction is complete c. equivalence point: volume of titrant added in excess of the end point to change a physical property of the analyte solution d. blank titration: titration performed without analyte to calculate the titration error e. standardization: titration of a known amount of analyte to determine the concentration of the titrant

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Chap 07_10e 5. Calculate pPb2+ when 25.00 mL of 0.100 M

is titrated with 21.00 mL of 0.100 M Pb2+. Ksp = 7.4 ×

10−14 for PbCO3. a. 11.07 b. 2.06 c. 11.41 d. 1.72 e. 11.99 6. Which of these practices help(s) protect the purity of chemical reagents?

I. Avoid putting a spatula into a bottle II. Never pour unused chemical back into the reagent bottle III. Never put a glass stopper from a liquid-reagent container on the lab bench IV. Store chemicals in a cool, dark place a. II b. II and III c. III and IV d. I, II, and III e. I, II, III, and IV 7. A 1.945 0-g sample containing only Na2CO3 (FM 105.99) and NaHCO3 (FM 84.01) is titrated with 35.31 mL of 0.872 4 M HCl. Calculate the percentage of sodium carbonate and sodium bicarbonate in the sample.

a. 43.53% Na2CO3; 56.47% Na2HCO3 b. 44.45% Na2CO3; 55.55% Na2HCO3 c. 55.55% Na2CO3; 44.45% Na2HCO3 d. 71.62% Na2CO3; 28.38% Na2HCO3 e. 56.47% Na2CO3; 43.53% Na2HCO3

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Chap 07_10e 8. For the precipitation titration of analyte X− by titrant M+, which of the statements is NOT true? a. The concentration of M+ at the equivalence point is equal to the molar solubility of MX, [M+] = . b. The concentration of M+ past the equivalence point is determined from the excess M+ after reacting with all of the X−. c. The concentration of M+ before the equivalence point is the allowed M+ for the [X−] remaining in the solution, −

.

d. The plot of the titration curve is [M+] versus VM+. e. The titration reaction is M+(aq) + X−(aq) → MX(s).

9. The end point for the Volhard titration is indicated by a. adsorption of a colored indicator on the precipitate at the end point. b. reduction of a colored indicator at the end point. c. deprotonation of a colored acid-base indicator at the end point. d. formation of a soluble, colored complex at the end point. e. protonation of a soluble, colored complex at the end point.

10. For the graphical equivalence point determination for a precipitation reaction, which statement(s) is/are TRUE? I. The steepest slope of the titration curve is the equivalence point.

II. The equivalence point occurs where the first derivative of the titration curve reaches its largest value.

III. The equivalence point occurs where the second derivative of the titration curve equals zero. a. I b. II and III c. II d. III e. I, II, and III 11. A compound available in high purity that does not decompose under ordinary storage conditions and is stable when dried by heating or vacuum is a a. secondary standard. b. primary standard. c. standard solution. d. indicator. e. titrant.

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Chap 07_10e 12. Calculate pCa2+ when 25.00 mL of 0.100 M

is titrated with 28.00 mL of 0.100 M Ca2+. Ksp = 1.3

× 10−8 for CaC2O4. a. 7.89 b. 2.58 c. 3.94 d. 2.25 e. 0.04 13. Calculate pBa2+ when 25.00 mL of 0.100 M

is titrated with 25.00 mL of 0.100 M Ba2+. Ks p = 1.1 ×

10−10 for BaSO4. a. 9.96 b. 7.14 c. 4.98 d. 1.30 e. 1.00 14. A Na2CO3 (FM 105.99) standard solution is prepared by transferring 2.481 7 g of primary standard-grade sodium carbonate to a 250.0-mL volumetric flask, dissolving the sample in ~100 mL of distilled deionized water, and diluting to the mark. A 25.00-mL aliquot is taken and titrated with 42.65 mL of an HCl solution.

Calculate the concentration of the HCl solution. a. 0.054 90 M b. 0.159 8 M c. 0.319 6 M d. 0.109 8 M e. 0.199 7 M 15. Argentometric titrations are titrations using a. Au+. b. Ag+. c. Ar+. d. Al3+. e. As3−.

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Chap 07_10e 16. Common___________ indicators are anionic dyes that are attracted to positively charged particles produced immediately after the equivalence point. a. redox b. absorption c. acid-base d. adsorption e. complexometric 17. The key step to titration calculations is to relate a. moles of titrant to moles of analyte. b. volume of titrant to volume of analyte. c. grams of titrant to grams of analyte. d. conductivity of titrant to conductivity of analyte. e. reduction of titrant to oxidation of analyte.

18. The Volhard method cannot be employed to analyze a. Cl−. b. SCN−. c. BH4−. d. . e. Zn2+.

19. A(n) ______________ is a compound with a physical property that changes abruptly near the equivalence point. a. primary standard b. titrant c. analyte d. indicator e. masking agent 20. The Fajans method cannot be employed to analyze a. Cl−. b. SCN−. c. BH4−. d. . e. Zn2+.

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Chap 07_10e 21. Calculate pPb2+ when 50.00 mL of 0.120 M I− is titrated with 25.00 mL of 0.125 M Pb2+. Ks p = 7.9 × 10−9 for PbI2.

22. Calculate pPb2+ when 25.00 mL of 0.150 M

is titrated with 17.00 mL of 0.150 M Pb2+. Ks p = 7.4 ×

10−14 for PbCO3.

23. Calculate pPb2+ at the equivalence point when 25.00 mL of 0.120 M

is titrated with 0.120 M Pb2+.

Ks p = 7.4 × 10−14 for PbCO3.

24. The amount of ascorbic acid (FM 176.12) in a sample can be determined by iodometric titration. A 10.00mL aliquot of a solution with an unknown ascorbic acid concentration was first pipetted into an Erlenmeyer flask, along with 5 mL of 5% KI solution and 1.5 mL of 6 M H2SO4. The mixture was then titrated rapidly with 1.897 mM KIO3 solution while swirling the flask between additions, until the yellow color of the complex almost persisted. Then four drops of starch indicator were added and the titration continued until the blue color of the starch-iodine complex persisted. The total amount of KIO3 solution used was 38.72 mL. Report the ascorbic acid concentration in mg ascorbic acid/100 mL solution.

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Chap 07_10e 25. Calculate pPb2+ when 25.00 mL of 0.200 M

is titrated with 28.00 mL of 0.200 M Pb2+. Ks p = 7.4 ×

10−14 for PbCO3.

26. Calculate pMg2+ when 50.00 mL of 0.120 M F− is titrated with 20.00 mL of 0.150 M Mg2+. Ks p = 7.4 × 10−9 for MgF2.

27. Calculate pAg+ when 25.00 mL of 0.152 M

is titrated with 32.00 mL of 0.145 M Ag+. Ks p = 1.2 ×

10−12 for Ag2CrO4.

28. Cl− can be determined by the Volhard method. A 25.00-mL aliquot of a Cl− solution of unknown concentration was treated with 2.30 mL of 6.0 M HNO3, and then precipitated by 30.00 mL of 0.486 2 M AgNO3 with vigorous stirring. Following the filtration of AgCl, the solution was mixed with 1.00 mL of saturated ferric alum, and titrated with 0.307 5 M KSCN. The solution turned red when 12.33 mL of KSCN solution had been added. What is the concentration Cl− in the unknown sample?

29. A NaOH solution is standardized using the monoprotic primary standard potassium hydrogen phthalate, KHP (FM 204.22). If 0.698 6 g of KHP requires 43.92 mL of NaOH, what is the NaOH concentration?

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Chap 07_10e 30. An aqueous mixture that is 0.100 M Ni2+ and 0.100 M Co2+ is titrated with a 0.100 M

solution.

Which cation precipitates first, and what is at the first and second equivalence points? Ks p = 1.0 × 10−10 for CoCO3 and Ks p = 1.3 × 10−7 for NiCO3.

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Chap 07_10e Answer Key 1. e 2. e 3. e 4. c 5. a 6. a 7. e 8. d 9. d 10. e 11. b 12. d 13. c 14. d 15. b 16. d 17. a 18. e 19. d 20. c 21. 2.78 22. 11.59 23. 6.57 24. 388.1 mg ascorbic acid/100 mL solution 25. 1.95 26. 2.91 Copyright Macmillan Learning. Powered by Cognero.

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Chap 07_10e 27. 5.17 28. 0.431 9 M 29. 0.077 89 M NaOH 30. Co2+ precipitates first;

= 5.00 at the first equivalence point; Ni2+ precipitates second;

= 3.44 at the

second equivalence point

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Chap 08_10e Indicate the answer choice that best completes the statement or answers the question.

1. What is the ionic strength of a 2:1 electrolyte with a concentration of 0.100 M? a. 0.100 b. 0.200 c. 0.300 d. 0.400 e. 0.500 2. Ionic ________ is a measure of the total concentration of ions in solution. a. radius b. density c. degree d. strength e. atmosphere 3. Which equation is NOT required to determine the pH of 0.10 M solution of the weak acid HA using the systematic treatment of equilibrium? a. HA(aq) ⇌ A−(aq) + H+(aq) Ka = 2.4 × 10−6

b. H2O(l) ⇌ H+(aq) + OH−(aq) Kw = 1.00 × 10−14 c. 0.10 M = [HA] + [A−] d. [H+] = [OH−] + [A−]

e. A−(aq) + H2O(l) ⇌ HA(aq) + OH−(aq)

Kb = 4.2 × 10−9

4. Which of these is NOT true for pH? a. pH = −log[H+]γH+. b. pH = −log[H+] when μ = 0. c. The pH = pOH for pure water. d. As ionic strength increases, so does the pH of water. e. The ionic strength of pure water is 1.0 × 10−7 M. Copyright Macmillan Learning. Powered by Cognero.

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Chap 08_10e 5. Which statement describes a mass balance for a solution prepared by dissolving 0.100 mol Na2SO4, 0.150 mol MgCl2, and 0.200 mol NaCl and diluting to 100.0 mL? a. The total Na+ concentration is 3.00 M. b. The total SO42– concentration is 0.100 M. c. The total Cl– concentration is 3.50 M. d. The total Mg2+ concentration is 15.0 M. e. None of the statements describe a mass balance for the solution.

6. For Ag+ when μ = 0.01 M, γ = 0.898 and when μ = 0.05 M, γ = 0.80. What is the activity coefficient when μ = 0.024 M? a. 0.804 b. 0.845 c. 0.853 d. 0.864 e. 0.894 7. What is the pH of a 0.001 0 M HNO3 solution with an ionic strength of 0.05 M? At μ = 0.05 M, γH+ = 0.86, γOH− = 0.81. a. 1.30 b. 1.37 c. 3.00 d. 3.07 e. 3.09 8. Calculate the pH of a 0.010 M CH3CO2H solution using the systematic treatment of equilibrium. Ignore activity coefficients. Ka = 1.8 × 10−5 for acetic acid. a. 9.26 b. 3.39 c. 2.02 d. 10.61 e. 4.74

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Chap 08_10e 9. Write the mass balance for a K2CO3 solution if the species in the solution are K+, CO32 –, HCO3 –, H2CO3, H+, and OH –. a. [K+] + [H+] = 2[CO32 –] + [HCO3 –] + [OH –] b. 2[K+] = [CO32 –] + [HCO3 –] + [H2CO3] c. [K+] = 2([CO32 –] + [HCO3 –] + [H2CO3]) d. [K+] = 2[CO32 –] + [HCO3 –] + [H2CO3] e. [K+] + [H+] = [CO32 –] + [HCO3 –] + [H2CO3] + [OH –] 10. Identify the mass balance equation for a 0.10 M H2C2O4 solution. a. [H+] = 2[C2O42 –] + [HC2O4 – ] + [OH – ] b. [H+] = 2[C2O42 –] + [HC2O4 – ] c. 0.10 M = 2[C2O42 –] + [HC2O4 – ] + [OH – ] d. 0.10 M = [C2O42 –] + [HC2O4 – ] + [H2C2O4] e. 2[H+] = [C2O42 –] + [HC2O4 –] + [H2C2O4] 11. Calculate the pOH of a 0.010 M HCl solution with an ionic strength of 0.10 M. At = 0.10 M , γH+ = 0.83, γOH– = 0.76. a. 11.92 b. 11.80 c. 11.88 d. 12.00 e. 12.20 12. Which statement is NOT correct for the pH of a 0.01 M NaCl solution versus the pH of a 0.01 M in FeSO4 solution? a. The activity coefficient for H+ is larger in 0.01 M NaCl than in 0.01 M FeSO4. b. The activity coefficient for OH− is larger in 0.01 M NaCl than in 0.01 M FeSO4. c. The H+ concentration in 0.01 M NaCl is equal to the H+ concentration in 0.01 M FeSO4. d. The H+ activity in 0.01 M NaCl is not equal to the H+ activity in 0.01 M FeSO4. e. The OH− activity in 0.01 M NaCl is not equal to the OH− activity in 0.01 M FeSO4.

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Chap 08_10e 13. A buffer is prepared from NaH2PO4 and Na2HPO4. Identify the charge balance equation for the buffer. a. [Na+] + [H+] = [H2PO4– ] + 2[HPO42–] + 3[PO43–] + [OH– ] b. 3[Na+] + [H+] = [H2 PO4– ] + 2[HPO42–] + 3[PO43–] + [OH– ] c. 3[Na+] + [H+] = [H2PO4– ] + [HPO42–] + [PO43–] + [OH– ] d. [Na+] = [H2PO4– ] + 2[HPO42–] + 3[PO43–] e. [Na+] + [H+] = [H2PO4– ] + [HPO42–] + [PO43–] + [OH– ] 14. Which equation is NOT required to determine the molar solubility of AgCN using the systematic treatment of equilibrium? a. AgCN(s) ⇌ Ag+(aq) + CN−(aq)

Ks p = 2.2 × 10−16

b. H2O(l) ⇌ H+(aq) + OH−(aq)

Kw = 1.00 × 10−14

c. [Ag+] = [CN−] + [HCN]

d. CN−(aq) + H2O(l) ⇌ HCN(aq) + OH−(aq)

Kb = 1.6 × 10−5

e. HCN(aq) ⇌ CN−(aq) + H+(aq)

Ka = 6.2 × 10−10

15. What is the ionic strength of 0.050 M Al(NO3)3? a. 0.050 b. 0.150 c. 0.200 d. 0.250 e. 0.300

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Chap 08_10e 16. Identify the charge balance equation for a Na2S solution if the species in the solution are Na+, S2 –, HS – , –

H2S, H+, and OH . a. [Na+] + [H+] = [S2 –] + [HS – ] + [OH – ] b. [Na+] + [H+] = 2[S2 –] + [HS – ] + [OH – ] c. 2[Na+] + [H+] = [S2 –] + [HS –] + [H2S] + [OH – ] d. 2[Na+] + [H+] = 2[S2 – ] + [HS – ] + [OH – ] e. [Na+] + [H+] = 2[S2 – ] + [HS4 – ] + [H2S] + [OH – ] 17. Which statements regarding the mass balance and charge balance of a solution are INCORRECT? I. Activity coefficients are used in the mass balance and the charge balance.

II. The mass balance is satisfied when the quantity of all species in a solution containing a particular atom (or group of atoms) equals the amount of that atom (or group) delivered.

III. The sum of the positive charges in solution equals the sum of the negative charges in solution.

IV. The coefficient in front of each species in a charge balance always equals the magnitude of the charge on the counterion. a. I, II, III, and IV b. I, III, and IV c. I, II, and III d. I and IV e. II and III 18. Which statement about activities and activity coefficients is FALSE? a. Activity for a chemical species is the product of its concentration and its activity coefficient. b. The activity coefficient corrects for nonideal behavior due to ionic strength. c. As ionic strength increases, the value of the activity coefficient increases. d. For ions, the activity coefficient approaches unity as the ionic strength approaches 0. e. The activity coefficient for neutral molecules is approximately unity when the ionic strength is less than 0.1 M.

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Chap 08_10e 19. Which statement is INCORRECT regarding the increased solubility of the sparingly soluble salt AgCl in an inert salt solution containing NaNO3? a. The ionic atmosphere around the silver cations becomes increasingly negative as the nitrate anions are attracted to the silver cations. b. The ionic strength around the chloride anions becomes increasingly positive as the sodium cations are attracted to the chloride anions. c. The ionic atmosphere increases the attraction between silver cations and chloride anions, increasing solubility by creating more AgCl(aq) in solution. d. Ions continually diffuse into and out of the ionic atmosphere. e. The negative ionic atmosphere around silver cations results in an overall positive charge less than the positive charge of the silver cations.

20. To correct for the effect of ionic strength on chemical reactions, concentrations are replaced by ________ in equilibrium expressions. a. activity coefficients b. ion activities c. molecular activities d. activities e. active strength 21. Which statement about activity coefficients is TRUE? a. As ion size increases, the value of the activity coefficient decreases. b. As the magnitude of the charge of the ion increases, the value of the activity coefficient decreases. c. As ionic strength increases, the value of the activity coefficient increases. d. The activity coefficient of a nonionic species is close to 0. e. None of above.

22. Use the systematic treatment of equilibrium to calculate the pH of a 0.002 0 M (CH3)3N solution. Ignore activity coefficients. Kb = 6.6 × 10−5 for (CH3)3N.

23. Use the systematic treatment of equilibrium to calculate the pH of a 50.0 mL solution containing 0.010 0 M HCl and 0.010 0 M formic acid. Ignore activity coefficients. Ka for formic acid is 1.80 × 10−4.

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Chap 08_10e 24. Use the systematic treatment of equilibrium to calculate the pH of a 0.001 0 M HCN solution. Ignore activity coefficients. Ka = 6.2 × 10−10 for HCN.

25. A Britton-Robinson buffer solution, a mixture of acetic acid, boric acid, and phosphoric acid, each at 0.040 M, was prepared and the pH was adjusted to 7.00 with 0.20 M NaOH. Calculate the ionic strength of the Britton-Robinson buffer solution at pH 7.00. The pKa values are as follows: acetic acid, 4.76; phosphoric acid, 2.15, 7.20, 12.33; and boric acid, 9.23.

26. Using activities, find [Ag+] in 0.010 M NaCl saturated with AgCl(s). Ks p = 1.8 10–10 for AgCl.

27. Calculate the ionic strength of a 0.210 M FeCl2 solution.

28. Calculate the pH of water containing 0.033 M BaCl2 at 25°C.

29. Calculate the activity coefficient of NO3 – in a 0.085 M NaNO3 aqueous solution at 25°C. The size of the nitrate ion is 300 pm.

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Chap 08_10e 30. Many metal cations, such as Al3+, are Lewis acids, which can react with water to produce hydrogen ions.

They may also precipitate as insoluble hydroxides. Write the charge balance of the solution, and the mass balance for Al, given the following equilibria in the solution.

Al3+ + H2O ⇌ Al(OH)2+ + H+ Al(OH)2+ + H2O ⇌ Al(OH)2+ + H+ Al(OH)2++ H2O ⇌ Al(OH)3 + H+ Al(OH)3 + H2O ⇌ Al(OH)4- + H+ H O ⇌ H+ + OH2

31. Calculate the activity coefficient for Fe2+ for an ionic strength of 0.034 M. For Fe2+, when m = 0.01 M, g = 0.675 and when m = 0.05 M, g = 0.485.

32. Calculate the pH of a 0.005 0 M NaOH solution that is also 0.005 0 M in KCl.

33. Calculate the pH of a 0.040 M HCl solution that is also 0.010 M in NaNO3.

34. Use the systematic treatment of equilibrium to calculate the molar solubility of TlIO3 in an aqueous solution with a pH of 2.00. Ignore activity coefficients. Ks p for TlIO3 is 3.1 × 10−6, and Ka for HIO3 is 0.17.

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Chap 08_10e Answer Key 1. c 2. d 3. c 4. d 5. e 6. d 7. d 8. b 9. c 10. d 11. a 12. c 13. a 14. e 15. e 16. b 17. d 18. c 19. c 20. d 21. b 22. pH = 10.52; Students can make approximations as they would for a simple weak-base equilibria problem and arrive at the same answer.

23. pH = 1.99; Students will need to list the charge balance, mass balance for formic acid, mass balance for the chloride ion, the ion product equilibrium expression, and the formic acid dissociation equilibrium expression to determine the hydronium ion concentration in the acid mixture. Students can neglect the concentration of the hydroxide ion in the charge balance due to the nature of an acidic solution.

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Chap 08_10e 24. pH = 6.10; Students can make approximations as they would for a simple weak-acid equilibria problem and arrive at the same answer.

25.

26. [Ag+] = 2.2 10–8 M; Determine the ionic strength of the solution and look up the activity coefficients for Ag+ and Cl– from Table 8-1.

27. 0.630; A straightforward ionic strength calculation.

28. pH = 6.98; Determine the ionic strength of the solution and look up the activity coefficients for H+ and OH– in Table 8-1. Use the activity coefficients to determine the H+ concentration in the solution from the Kw expression. The pH of the solution is determined using activity.

29. γ = 0.749; First, determine the ionic strength of the solution. Then apply the extended Debye-Hückel equation.

30. Charge balance: 3[Al3+] + 2[Al(OH)2+] + [Al(OH)2+] + [H+] = [Al(OH)4-] + [OH-] Mass balance for total Al = [Al3+] + [Al(OH)2+] + [Al(OH)2+] + [Al(OH)3] + [Al(OH)4-] 31. g = 0.561 for Fe2+ when m = 0.034 M; A straightforward interpolation problem.

32. pH = 11.65; This problem requires Table 8-1 for the students to look up the activity coefficients for OH– . gOH– = 0.900 when m = 0.010 M. Since NaOH is a strong base, the hydroxide concentration is 0.005 0 M, and the activity of hydroxide is 0.004 5 M. Therefore, the pH of the solution is 11.65.

33. pH = 1.46; Since HCl is a strong acid, the H+ concentration from the dissociation of HCl is 0.040 M. The ionic strength of this solution is 0.050 M. This problem requires Table 8.1 for the students to look up the activity coefficients for H+. gH+ = 0.86 when m = 0.050 M. Calculate the pH using the activity coefficient and concentration of H+ (pH = –log gH+·[H+]) 34. molar solubility = 1.8 × 10−3 M; Students will need to list the solubility product equilibrium expression, the mass balance for the solution, the ion product equilibrium expression, and the iodate hydrolysis equilibrium expression to determine the Tl+ concentration, which is equal to the molar solubility. No charge balance is required because the solution is buffered at pH 2.00.

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Chap 09_10e Indicate the answer choice that best completes the statement or answers the question.

1. Given the formal concentration and Ka of a weak acid (HA), the pH of the weak acid solution can be calculated using the systematic treatment of equilibrium. Which equation is NOT needed to solve for the pH of the solution? a. b. F = [HA] + [A−] c. [H+] = [A−] + [OH−] d. e. Kw = [H+][OH−] 2. A student calculates the pOH for a 0.010 M triethanolamine (Kb = 5.78 × 10−7) solution to be 4.12. To simplify the math, she decided to ignore the change in the formal concentration in her calculation. Was she correct in her decision? a. She was correct to ignore the change, as the change in the formal concentration, x, was 7.59 × 10−5, which is less than 1% of the formal concentration. b. She was not correct to ignore the change, as the change in the formal concentration, x, was 9.92 × 10−3, which is greater than 1% of the formal concentration. c. She was correct to ignore the change, as the change in the formal concentration, x, was 1.32 × 10−10, which is less than 1% of the formal concentration. d. She was not correct to ignore the change, as the change in the formal concentration, x, was 1.00 × 10−2, which is greater than 1% of the formal concentration. e. She was correct to ignore the change, as the change in the formal concentration, x, was 5.78 × 10−5, which is less than 1% of the formal concentration.

3. Which statement is NOT true for weak bases (B)? a. Weak bases are partially hydrolyzed in aqueous solution. b. The conjugate acid is BH+ and it is a strong acid. c. d. For the conjugate acid,

.

e. B(aq) + H2O(l) ⇌ BH+(aq) + OH−(aq)

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Chap 09_10e 4. Calculate the pH of 0.002 5 M NaOH if activities are neglected. a. −2.60 b. 2.60 c. 7.00 d. −11.40 e. 11.40 5. Calculate the pH of 0.100 M NaCN. The Ka for HCN is 9.12 × 10−10. a. 2.98 b. 11.02 c. 4.52 d. 9.48 e. 9.04 6. A 0.010 M solution of a weak base has a percent association of 18.32%. Calculate the Kb for the weak base. a. 1.83 × 10−3 b. 2.43 × 10−11 c. 5.46 × 10−12 d. 4.11 × 10−4 e. 1.22 × 10−12 7. A buffer is prepared by mixing 200.0 mL of 0.150 0 M NaOH with 200.0 mL of 0.200 0 M CH3CO2H (Ka = 1.75 × 10−5) in a 1-L volumetric flask, and then diluting to volume with deionized water. Calculate the pH of the buffer. a. 5.23 b. 9.72 c. 7.00 d. 8.77 e. 4.28

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Chap 09_10e 8. There are six common strong acids. Which statement is NOT true for strong acids? a. Strong acids are completely dissociated in dilute solutions. b. HNO2 is a strong acid. c. For strong acid solutions with a concentration between 10−6 and 10−8 M, the pH of the solution is determined using the systematic treatment of equilibrium. d. For strong acid solutions with concentrations ≥ 10−6 M, the pH of the solution is calculated from the concentration of the strong acid. e. For strong acid solutions with concentrations ≤ 10−8 M, the pH of the solution is always 7.

9. In a 0.000 500 M HBr solution, the dissociation of water produces ___________ H3O+ and ___________ OH−. a. 1.00 × 10−7 M; 2.00 × 10−11 M b. 5.00 × 10−4 M; 5.00 × 10−4 M c. 2.00 × 10−11 M; 2.00 × 10−11 M d. 1.00 × 10−7 M; 1.00 × 10−7 M e. 5.00 × 10−4 M; 2.00 × 10−11 M 10. Given the formal concentration and Kb of a weak base (B), the pH of the weak base solution can be calculated using the systematic treatment of equilibrium. Which equation is NOT needed to solve for the pH of the solution? a. b. F = [BH+] + [B] c. [OH−] = [BH+] + [H+] d. e. Kw = [H+][OH−] 11. A buffer is made of a weak acid (HA) and its conjugate base (A−). When [HA]:[A−] is 10:1, the pH of the buffer is equal to a. pKa + 10. b. pKa + 1. c. pKa. d. pKa – 1. e. pKa – 10.

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Chap 09_10e 12. Calculate the percent dissociation for 0.010 M hypochlorous acid. The Ka for hypochlorous acid is 3.0 × 10−8. a. 0.35% b. 0.30% c. 0.17% d. 0.99% e. 0.65% 13. Which of the following is NOT true for weak acids (HA) and weak bases (B)? a. Kw = Ka·Kb b.

c. d. pKa = −log Ka e. pKb = −log Kb 14. Strong bases consist of alkali metal hydroxides and alkaline earth metal hydroxides, although the latter are far less soluble. Which statement is NOT true for strong bases? a. Strong bases are completely dissociated in dilute solutions. b. Fe(OH)2 is a strong base. c. For strong base solutions with a concentration between 10−6 and 10−8 M, the pOH of the solution is determined using the systematic treatment of equilibrium. d. For strong base solutions with concentrations ≥ 10−6 M, the pOH of the solution is calculated from the concentration of the strong base. e. For strong base solutions with concentrations ≤ 10−8 M, the pOH of the solution is always 7.

15. In an equilibrium problem involving a weak acid or a weak base, the change in the formal concentration of the weak acid or the weak base, x, can be ignored if x is a. greater than 1% of the formal concentration. b. less than 0.1% of the formal concentration. c. equal to 0.1% of the formal concentration. d. less than 10% of the formal concentration. e. less than 1% of the formal concentration.

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Chap 09_10e 16. Which is NOT a property of buffers? a. Buffers resist changes in pH when a small amount of acid or base is added to the buffer or when a small dilution of the buffer occurs. b. Buffers are a mixture of a weak acid and its conjugate base. c. The pH of a buffer is independent of ionic strength. d. The pH of a buffer is dependent on temperature. e. Buffers are a mixture of a weak base and its conjugate acid.

17. Formic acid (HCO2H) is a weak acid with a Ka value of 1.80 × 10−4. What is the Kb value of its conjugate base, formate ( )? a. 1.80 × 10−4 b. 5.56 × 10−11 c. 1.80 × 1010 d. 1.80 × 10−14 e. 5.56 × 1011 18. Which statement is NOT true for weak acids (HA)? a. Weak acids are partially dissociated in aqueous solution. b. The conjugate base is A− and it is a weak base. c. d. For the conjugate base,

.

e. HA(aq) + H2O(l) ⇌ A−(aq) + H3O+(aq) 19. Calculate the pH of 0.005 0 M HF. The Ka for HF is 6.8 × 10−4. a. 1.58 b. 2.73 c. 0.43 d. 3.55 e. 2.81

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Chap 09_10e 20. The buffer required for a capillary electrophoresis experiment must have a pH of 10.0. Which weak acid is the best choice for the buffer? Assume = 0. a. ammonium chloride pKa = 9.24 b. disodium phosphate

pKa = 12.38

c. 3-(cyclohexylamino)propanesulfonic acid

pKa = 10.40

d. sodium dihydrogen borate

pKa = 12.74

e. cyclohexylaminoethanesulfonic acid

pKa = 9.39

21. Calculate the pH of 0.005 0 M phenylacetic acid. The Ka for phenylacetic acid is 4.90 × 10–11.

22. Calculate the percent dissociation for a 0.010 0 M HCN solution. Ka = 6.2 × 10−10 for HCN.

23. You need to prepare 1.00 L of a pH 5.00 propionic acid buffer with a total concentration of 50 mM. Calculate the millimoles of propionic acid and sodium propionate needed to prepare the buffer. Ka = 1.34 × 10−5 for propionic acid.

24. Including activity coefficients, calculate the pH of a 4.3 × 10−8 M HCl solution with an ionic strength of m = 0.10. The activity coefficients are , , and .

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Chap 09_10e 25. Calculate the pH of a 0.200 M ethylamine solution. Kb = 5.0 × 10−4 for ethylamine.

26. Calculate the pH of 0.010 M HCN (m = 0.10 M) using activity coefficients and neglecting activity coefficients. Then, determine the percent difference between the two values. Ka = 6.2 × 10−10 for HCN. The activity coefficients are , , and .

27. Calculate the pH of 0.050 0 M NaF. The Ka for HF is 6.8 × 10−4.

28. Chlorous acid (HClO2) is a weak acid with a Ka value of 1.10 × 10−2. What is the Kb value of its conjugate base, chlorite

?

29. How many milliliters of 0.105 M HCl should be added to 200.0 mL of 0.010 9 M triethylamine to give a buffer with a pH of 10.20? The Kb of triethylamine is 5.26 × 10−4.

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Chap 09_10e 30. Including activity coefficients, calculate the pH of 0.005 0 M KOH. Ion m = 0.001 M m = 0.005 M m = 0.01 M 0.967 0.933 0.914 H+ 0.964 0.926 0.900 OH− 0.964 0.925 0.899 K+

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m = 0.05 M 0.86 0.81 0.805

m = 0.1 M 0.83 0.76 0.755

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Chap 09_10e Answer Key 1. d 2. a 3. b 4. c 5. b 6. d 7. a 8. b 9. c 10. a 11. d 12. c 13. b 14. b 15. e 16. c 17. b 18. d 19. e 20. c 21. pH = 6.31; Solve the weak-acid equilibrium problem, ignoring the change in the formal concentration of phenylacetic acid.

22. 0.0249%; Solve the weak-acid equilibrium for [CN−], ignoring the change in the formal concentration of HCN. Divide the concentration of the cyanide ion by the formal concentration and multiply by 100.

23. 28.6 mmol sodium propionate and 21.4 mmol propionic acid; Set millimoles of propionate equal to x and millimoles of propionic acid equal to 50 − x. Insert into the Henderson-Hasselbalch equation and solve for x.

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Chap 09_10e 24. pH = 6.59; The concentration of HCl is less than the hydronium concentration from the autoprotolysis of water, so [OH−] is calculated from Kw and the concentration of HCl, keeping in mind that for each OH− in water, a H+ is also generated. Add [H+] from Kw and from HCl to get total [H+]. Use activity coefficients where needed and when calculating the pH. When activities are neglected, pH = 6.69.

25. pH = 12.30; Solve the weak-base equilibrium problem for [OH−] without assuming that the change in the ethylamine formal concentration is less than 1%. If students report pH = 12.00, then they failed to notice that neglecting the change in the formal concentration results in an error of 5%.

26. With activity coefficients pH = 5.20; without activity coefficients pH = 5.30. The pH calculated with activity coefficients is 1.92% less than the pH calculated without activity coefficients. Solve the weak-acid equilibrium problem with and without activity coefficients. Then calculate the percent difference between the pH values. If the pH value calculated without activity coefficients is used in the denominator when calculating the percent difference, the difference is 1.89%.

27. 7.93; Determine Kb for F− from the Ka for HF. Solve the weak-base equilibrium problem for [OH−], ignoring the change in the formal concentration of F−. Calculate the pOH and then the pH. 28. 9.09 × 10–13 29. 15.9 mL; Calculate the millimoles of triethylamine. Set the millimoles of the triethylammonium ion equal to x and the millimoles of triethylamine equal to the initial number of moles minus x. Calculate pKa from the Kb , insert the values into the Henderson-Hasselbalch equation, and solve for x, which is equal to the millimoles of HCl needed. Use the concentration of the HCl solution to calculate the volume.

30. pH = 11.67; Determine the ionic strength of the solution. Then use the concentration of KOH and the appropriate activity coefficient to calculate pOH. Then calculate the pH. When activities are neglected pH = 11.70.

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Chap 10_10e Indicate the answer choice that best completes the statement or answers the question.

1. For a monoprotic weak acid, αHA ____________ and αA – ________ as the pH increases. a. increases; decreases b. increases; increases c. decreases; decreases d. decreases; increases e. remains constant; remains constant 2. How many millimoles of NaOH or HCl must be added to 500.0 mL of 0.250 M NaO2CCO2H to prepare a pH 3.90 buffer? Assume no change in volume. pK1 = 1.250 and pK2 = 4.266. a. 37.65 mmol NaOH b. 87.35 mmol HCl c. 124.7 mmol HCl d. 0.30 mmol NaOH e. The answer cannot be calculated from the given information.

3. Calculate the pH of a 0.340 M NaO2CCO2H solution. pK1 = 1.250 and pK2 = 4.266 for HO2CCO2H. a. 2.37 b. 4.89 c. 2.79 d. 3.38 e. −0.029 4. Calculate the pH of a 0.100 M sodium alaninate solution. pKb1 = 4.29 and pKb2 = 11.67. a. 2.65 b. 11.35 c. 6.34 d. 7.66 e. 7.98

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Chap 10_10e 5. Sulfurous acid is a diprotic acid with Ka1 = 1.40 × 10−2 and Ka2 = 6.73 × 10−8. Which statements are TRUE? I. Kb2 = 7.14 × 10−13 II. Kb1 = 1.49 × 10−7 III. For a 0.100 M H2SO3 solution, [H+] = 7.61 × 10−2 M for the first dissociation.

IV. For a 0.100 M NaHSO3 solution, pH ≈ 4.513. a. I, II, and IV b. I, II, and III c. I and II d. I, II, III, and IV e. III and IV 6. For the polyprotic acid H3A, which equilibria and corresponding K value is INCORRECT? a. A3− + H2O ⇌ HA2− + OH−

Kb3 = Kw/Ka3

b. H2A− ⇌ HA2− + H+

Ka2 = K2

c. HA2− ⇌ A3− + H+

Ka3 = K3

d. H3A ⇌ H2A− + H+

Ka1 = K1

e. HA2− + H2O ⇌ H2A− + OH−

Kb2 = Kw/Ka2

7. Phenylalanine (F) is a diprotic amino acid with pK1 = 2.20 for the carboxylic acid group and pK2 = 9.31 for the ammonium group. At pH = 5.93, the principal species in solution is _______, and at pH = 10.10, the principal species in solution is _________. a. H2F+; H2F+ b. HF; F− c. H2F+; F− d. H2F+; HF e. HF; HF

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Chap 10_10e 8. Alanine (A) is a diprotic amino acid with pK1 = 2.344 for the carboxylic group and pK2 = 9.868 for the ammonium group. At pH = 6.32, what is the principal species in solution? a. H2A+ b. A− c. HA d. H3A2+ e. A2− 9. Isoleucine (I) is a diprotic amino acid with pK1 = 2.318 for the carboxylic acid group and pK2 = 9.758 for the ammonium group. What is TRUE when the pH of the solution is equal to pK1 for isoleucine? a. H2I+ is the predominant species. b. [H2I+] = [HI] c. HI is the predominant species. d. [HI] = [I−] e. I− is the predominant species.

10. Calculate the pH at the isoelectric point of a 0.100 M solution of the amino acid threonine. pK1 = 2.088 and pK2 = 9.100 for threonine. a. 3.57 b. 2.39 c. 5.59 d. 1.54 e. 5.05 11. Calculate the pH of a 0.200 M H2A− solution. K1 = 1.00 × 10−4, K2 = 1.00 × 10−8 and K3 = 1.00 × 10−12. a. 4.35 b. 6.00 c. 8.00 d. 8.65 e. 10.00

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Chap 10_10e 12. The fractional composition diagram is representative of a diprotic weak acid. What is the principal species at point E?

a. H2A b. HA− c. A2− d. H2A = HA− e. HA− = A2− 13. Nitrilotriacetic acid ( ) is a tetraprotic acid with pK1 = 1.0, pK2 = 1.81, pK3 = 2.52, and pK4 = 9.46. What is the principal species of nitrilotriacetic acid in a solution with pH = 6.50? a. b. C6H9NO6 c. d. e.

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Chap 10_10e 14. Calculate the pH of a buffer prepared by dissolving 0.840 g NaHCO3 (FM 84.007) and 0.530 g Na2CO3 (FM 105.989) in 250.0 mL water. pK1 = 6.13 and pK2 = 9.91. a. 8.02 b. 5.83 c. 6.43 d. 9.61 e. 10.21 15. Which statements are TRUE for an aqueous solution of the polyprotic acid H3A? I. H3A is acidic.

II. H2A− is amphiprotic.

III. HA2− is amphiprotic.

IV. A3− is basic. a. I and IV b. I, II, and IV c. II and III d. I, III, and IV e. I, II, III, and IV 16. Calculate the pH of a 0.012 8 M H2SO3 solution. pK1 = 1.857 and pK2 = 7.172 for sulfurous acid. a. 1.87 b. 2.09 c. 1.86 d. 1.89 e. 4.51 17. Citric acid is a triprotic acid (H3A) with K1 = 7.44 × 10−4, K2 = 1.73 × 10−5, and K3 = 4.02 × 10−7.

What is the Kb1 value for citrate (A3–)? a. 1.34 × 10−11 b. 5.78 × 10−10 c. 2.49 × 10−8 d. 7.44 × 10−4 e. 1.73 × 10−5

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Chap 10_10e 18. A buffer is prepared by adding 70 mmol of NaOH(s) to 1.00 L of a 50 mM H2A solution. Assume no change in volume of the solution. Which statements regarding the buffer are FALSE? I. There is 30 mmol HA− and 20 mmol A2− in the solution.

II. The pH is calculated using

.

III. There is 50 mmol HA− and 20 mmol NaOH in solution.

IV. pH is calculated using

.

V. pH = 14 – pOH, where pOH is calculated from excess OH−. a. III, IV, and V b. II, III, IV, and V c. II, III, and IV d. I, II, and V e. I and II 19. Which statement is NOT true of the isoelectric or isoionic pH? a. The isoionic pH is the pH of the pure, neutral, polyprotic acid. b. The isoelectric pH is the pH at which the average charge of the polyprotic acid is +1. c. The hydronium concentration at the isoionic point for a diprotic acid is calculated with the equation . d. The pH at the isoelectric point for a diprotic acid is calculated with the equation pH = ½(pK1 + pK2). e. In the equation

, F is the formal concentration of the neutral species of

the diprotic acid.

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Chap 10_10e 20. Which statement is INCORRECT for diprotic acids and bases? a. The diprotic species H2A can be treated as monoprotic to calculate [H+], [HA−], and [H2A] in an H2A solution. b. The pH for a solution containing the intermediate species HA− is approximated with pH = ½(pK1 + pK2). c. Ka1Kb1 = Kw and Ka2Kb2 = Kw d. The dibasic species A2− can be treated as monobasic to calculate [OH−], [HA−], and [A2−] in an A2− solution. e. A zwitterion is a neutral molecule with both positive and negative charges.

21. Calculate the pH of a 0.002 50 M hemimellitic acid solution. pK1 = 2.86, pK2 = 4.30, and pK3 = 6.28 for hemimellitic acid.

22. Calculate the pH of a solution prepared by mixing 200.0 mL of 0.100 M NaOH with 150.0 mL of 0.200 M oxalic acid. K1 = 5.62 × 10−2 and K2 = 7.24 × 10−5 for oxalic acid.

23. Calculate the concentration of each species of oxalic acid in a 0.150 M solution of sodium oxalate, Na2C2O4. K1 = 5.62 × 10−2 and K2 = 7.24 × 10−5 for oxalic acid.

24. Aspartic acid is a triprotic acid with pK1(H3D+) = 1.990, pK2(H2D) = 3.900, and pK3(HD−) = 10.002.

Calculate the pH for a 0.200 M H2D solution.

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Chap 10_10e 25. Calculate the pH of a 0.150 M maleic acid solution. pK1 = 1.92 and pK2 = 6.27 for maleic acid.

26. Calculate the pH of 0.016 6 M K3PO4 solution. For phosphoric acid (H3PO4), pK1 = 1.92, pK2 = 6.71, and pK3 = 11.52

27. Calculate the fraction of ethanolamine that is not pronated in a solution at pH 10.00. pKa for ethanolamine (HOCH2CH2NH3+) is 9.52.

28. Malonic acid is a diprotic acid with pK1 = 2.847 pK2 = 5.696. Calculate the fraction of malonic acid in each of its forms at pH = 3.50.

29. Calculate the pH of a 0.033 M Na2SO3 solution. pK1 = 1.66 and pK2 = 6.85 for sulfurous acid.

30. Calculate the pH to which a solution of the amino acid alanine must be adjusted to reach its isoelectric point. pK1 = 2.344 and pK2 = 9.868.

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Chap 10_10e Answer Key 1. d 2. a 3. c 4. b 5. c 6. a 7. b 8. c 9. b 10. c 11. b 12. e 13. d 14. d 15. e 16. b 17. c 18. a 19. b 20. c 21. 2.89; Treat hemimellitic acid as a monoprotic acid. Solve the equilibrium for [H+] using the quadratic equation. Do not assume that the change in the formal concentration is negligible.

22. 1.55; The mmol OH− added is sufficient to only convert some of the H2A to HA−, which gives a buffer centered on the pK1.

23. [A2−] = 0.150 M, [HA−] = 4.55 × 10−6 M, [H2A] = 1.78 × 10−13; Calculate Kb1 and Kb2 from the K values for oxalic acid. Treat the oxalate ion as a monobasic base, and solve the equilibrium for the concentrations of A2− and HA−. The concentration of H2A is equal to Kb2.

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Chap 10_10e

24. 2.96; The [H+] is calculated using the equation for the first intermediate form,

.

25. 1.43; Treat maleic acid as a monoprotic acid and solve the equilibrium for [H+] using the quadratic equation. Then calculate the pH. If you incorrectly ignore the change in the formal concentration of maleic acid, the pH is 1.37.

26. 11.77; Calculate Kb1 from K3 for phosphoric acid. Treat the phosphate ion as a monobasic base. Solve the equilibrium for [OH−] using the quadratic formula, and calculate the pOH and pH. Do not assume that the change in the formal concentration is negligible.

27. 0.751 28.

,

,

; Insert the H+ concentration and Ka values into the

alpha equations for H2A, HA−, and A2−.

29. 9.68; Determine Kb1 from K2. Treat Na2SO3 as monobasic and solve the equilibrium for [OH−], using the assumption that the change in the formal concentration is negligible. Then calculate pOH and pH. 30. 6.106; The pH is the average of the pKa values.

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Chap 11_10e Indicate the answer choice that best completes the statement or answers the question.

1. For the titration of a strong acid with a strong base, which statement is FALSE? a. At V = 0 mL titrant added, the pH is determined using the initial concentration of the strong acid. b. At the equivalence point, the pH is determined by the autoprotolysis of water. c. Between V = 0 mL and the equivalence point, the pH is determined using the diluted concentration of the unreacted strong acid. d. Past the equivalence point, the pH is determined using the diluted concentration of the excess strong base. e. The indicator changes color at the end point due to the presence of excess acid.

2. The end point for an acid-base titration can be determined graphically. The end point volume on a first derivative plot corresponds to the volume of titrant added where a. the first derivative curve crosses the x-axis. b. the first derivative curve reaches a maximum. c. the first derivative curve crosses the y-axis. d. the first derivative curve reaches a minimum. e. the slope of the first derivative curve is the greatest.

3. A 50.00-mL solution of 0.100 M HCl is titrated with 0.150 M NaOH. What is the pH at the equivalence point? a. 1.00 b. 13.18 c. 1.22 d. 12.78 e. 7.00 4. Which statement regarding the titration of a weak acid with a strong base is FALSE? HA + OH− → A− + H2O a. At the equivalence point, the pH is determined by the hydrolysis of the conjugate base, A−. b. Past the equivalence point the pH is largely determined by excess conjugate base, A−. c. Before any strong base is added, the pH is determined by the dissociation of the weak acid, HA. d. Before the equivalence point, the solution is a buffer, containing a mixture of HA and A−. e. The pH at the equivalence point is basic.

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Chap 11_10e 5. Arrange the following steps of the Kjeldahl nitrogen analysis procedure in order from first to last.

I. Distill ammonia into a standard hydrochloric acid solution.

II. Neutralize ammonium ion to release ammonia.

III. Digest an organic sample with boiling sulfuric acid to convert nitrogen to ammonium ion.

IV. Titrate unreacted hydrochloric acid with sodium hydroxide. a. III, II, I, IV b. II, I, IV, III c. I, IV, III, II d. IV, III, II, I e. III, I, IV, II 6. Which statement regarding indicators for acid-base titrations is FALSE? a. Indicators are acids or bases whose various protonated forms have different colors. b. The pH range over which an indicator changes color is called the transition range. c. The point at which the indicator changes color is the true equivalence point of the titration. d. The volume difference between the equivalence point and the end point is the indicator error. e. Indicators can change color more than once.

7. Which statement regarding the titration of a weak base with a strong acid is FALSE? B + H+ → BH+ a. At the equivalence point, the pH is determined by the dissociation of the conjugate acid, BH+. b. Past the equivalence point, the pH is determined from the excess strong acid, H+. c. Before strong acid is added, the pH is determined by the hydrolysis of the weak base, B. d. Before the equivalence point, the solution is a buffer, containing a mixture of B and BH+. e. The pH at the equivalence point is neutral.

8. Which statements regarding Gran plots are TRUE? I. A Gran plot is more accurate near the end point than is a first or second derivative plot.

II. A Gran plot is created using pH data at added titrant volumes just before the end point.

III. The end point of the titration is determined by extrapolating the linear portion of the Gran plot to where it crosses the x-axis.

IV. A Gran plot can be used with polyprotic titrations as well as monoprotic titrations. a. I, II, and IV b. II, III, and IV c. I, II, and III d. I, II, III, and IV e. I, III, and IV

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Chap 11_10e 9. Which statement regarding the titration of a dibasic base with a strong acid is FALSE? B + H+ → BH+ BH+ + H+ → a. At V = 0 titrant added, the pH is determined by the hydrolysis of B. b. Before the first equivalence point, the pH is determined by the buffer created from the mixture of B and BH+. c. At the first equivalence point, B has been converted to BH+. This is the most buffered point on the whole titration curve. d. Between the first and second equivalence points, the pH is determined by the buffer created from the mixture of BH+ and . e. At the second equivalence point, the pH is determined by the dissociation of

.

10. A strong-base solution requires standardization before use because strong bases ______________, which decreases the hydroxide concentration of the solution. a. precipitate as insoluble metal hydroxides b. react with atmospheric carbon dioxide to form carbonates c. react with atmospheric nitrogen to form nitrogen oxides d. adhere to the inside surface of the solution container e. react with solution impurities to form volatile compounds 11. A 25.00-mL aliquot of 0.090 0 M acetic acid is titrated with 0.100 M KOH. Calculate the pH after 25.00 mL of 0.100 M KOH has been added. Ka = 1.75 × 10−5 for acetic acid. a. 2.30 b. 5.27 c. 5.71 d. 8.73 e. 11.70 12. Which compound CANNOT be used as a primary standard in the titration of an acid or a base? a. potassium hydrogen phthalate b. sodium carbonate c. sodium hydroxide d. potassium hydrogen iodate e. benzoic acid

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Chap 11_10e 13. A 25.00-mL aliquot of a 0.100 M weak base solution is titrated with 0.100 M HCl. What is the approximate pH at the equivalence point? a. pH = 1.00 b. 1.00 < pH < 7.00 c. pH = 7.00 d. 7.00 < pH < 13.00 e. pH = 13.00 14. A 25.00-mL aliquot of a 0.100 M weak acid (Ka = 1.0 × 10−4) solution is titrated with a strong base. What is the approximate pH at the half-equivalence point? a. 2.00 b. 4.00 c. 6.00 d. 8.00 e. 10.00 15. The pH at the equivalence point of the titration of a weak acid with a strong base is 8.04. The pH increases steeply from 6.5 to 9.5 over a small volume interval. Which indicator would be the best for observing the end point of this titration? a. bromocresol purple – transition range 5.2–6.8 b. phenolphthalein – transition range 8.0–9.6 c. bromothymol blue – transition range 6.0–7.6 d. chlorophenol red – transition range 4.8–6.4 e. methyl orange – transition range 3.1–4.4 16. The end point for an acid-base titration can be determined graphically. The end point volume on a second derivative plot corresponds to the volume of titrant added where a. the second derivative curve crosses the x-axis. b. the second derivative curve reaches a maximum. c. the second derivative curve crosses the y-axis. d. the second derivative curve reaches a minimum. e. the slope of the second derivative curve is the greatest.

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Chap 11_10e 17. Which statements regarding the leveling effect are TRUE? I. Acids stronger than water react with water to form hydronium.

II. Hydronium is the strongest acid in water.

III. In water, strong acids are all leveled to the same strength.

IV. The differences in the strength of strong acids are observed in nonaqueous solvents such as acetic acid. a. I, II, and III b. I and II c. II and III d. I, II, III, and IV e. I and III 18. A 25.00-mL aliquot of 0.100 M Na2CO3 is titrated with 0.100 M HCl. What is the pH at the first equivalence point? pK1= 6.13 and pK2 = 9.91 for carbonic acid. a. 6.13 b. 9.91 c. 8.02 d. 7.00 e. 5.42 19. A 40.00-mL aliquot of 0.100 0 M malonic acid is titrated with 0.090 0 M KOH. What volume of 0.090 0 M KOH must be added to give a pH of 5.697? K1 = 1.42 × 10−3 and K2 = 2.01 × 10−6 for malonic acid. a. 66.67 mL b. 44.44 mL c. 22.22 mL d. 88.89 mL e. 111.11 mL 20. A 25.00-mL aliquot of 0.100 M NH3 is titrated with 0.090 0 M HCl. Calculate the pH after 21.00 mL of 0.090 0 M HCl has been added. Kb = 1.76 × 10−5 for ammonia. a. 5.25 b. 9.74 c. 8.75 d. 4.26 e. 7.00

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Chap 11_10e 21. Strong-base solutions should be stored in I. clear glass containers.

II. amber glass containers.

III. plastic containers. a. I b. II c. III d. I and II e. I, II, and III 22. A 25.00-mL aliquot of 0.100 M KOH is titrated with 0.065 2 M HCl. Calculate the pH after 40.00 mL of 0.065 2 M HCl has been added.

23. A 50.00 mL aliquot of 0.150 0 M methylamine (Kb = 4.42 × 10−4) is titrated with 0.100 0 M HCl. Calculate the pH after 35.00 mL of 0.100 0 M HCl has been added.

24. A 25.00-mL aliquot of 0.100 M sodium acetate (pKb = 9.244) is titrated with 0.125 M HCl. Calculate the pH after 10.00 mL of 0.125 M HCl has been added.

25. A 25.00-mL aliquot of 0.100 M oxalic acid (HO2CCO2H) is titrated with 0.100 M KOH. Calculate the pH after 35.00 mL of 0.100 M KOH has been added. K1 = 5.4 × 10−2 and K2 = 5.42 × 10−5 for oxalic acid.

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Chap 11_10e 26. A 50.00-mL aliquot of 0.100 0 M propenoic acid (H2CCHCO2H) is titrated with 0.125 0 M NaOH. Calculate the pH after 25.00 mL of 0.125 0 M NaOH has been added. Ka = 5.52 × 10−5 for propenoic acid.

27. A 25.00-mL mixture of NaHCO3 and Na2CO3 was titrated with 0.100 0 M HCl, and two equivalence points were observed at V = 17.53 mL and V = 46.28 mL. Determine the molarity of NaHCO3 and Na2CO3 in the original mixture.

28. A 1.000-g organic sample containing nitrogen is subjected to Kjeldahl analysis. The ammonia gas generated when the ammonium ion is neutralized is distilled into 50.00 mL of 0.100 0 M HCl to reform the ammonium ion. If 17.29 mL of 0.100 0 NaOH is required to titrate the unreacted HCl, calculate the percent by mass nitrogen in the organic sample.

29. A 25.00-mL aliquot of 0.100 M benzoic acid (pKa = 4.202) is titrated with 0.125 M NaOH. Calculate the pH after 20.00 mL of 0.125 M NaOH has been added.

30. A 50.00-mL aliquot of 0.100 M maleic acid (HO2C(CH)2CO2H) is titrated with 0.100 M NaOH. Calculate the pH after 20.00 mL of 0.100 M NaOH has been added. K1 = 1.20 × 10−2 and K2 = 5.37 × 10−7 for maleic acid.

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Chap 11_10e 31. A 50.00-mL aliquot of 0.100 M HNO3 is titrated with 0.075 0 M NaOH. Calculate the pH after 40.00 mL of 0.075 0 M NaOH has been added.

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Chap 11_10e Answer Key 1. e 2. b 3. e 4. b 5. a 6. c 7. e 8. d 9. c 10. b 11. e 12. c 13. b 14. b 15. b 16. a 17. d 18. c 19. a 20. c 21. c 22. 2.779; Determine the number of moles of HCl added and the number of moles of KOH in the solution. The difference is the moles of excess HCl. Determine the concentration of the excess HCl using the total volume. Then calculate the pH. 23. 10.70; At this point in the titration, a buffer exists with 3.5 mmol BH+ and 4.0 mmol B. Use the pKa of the methylammonium ion and the Henderson-Hasselbalch equation to calculate the pH. 24. 4.756; This is halfway to the equivalence point, so pH = pKa for acetic acid.

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Chap 11_10e 25. 4.090; The added base exceeds the first equivalence point, but is less than required to reach the second equivalence point. The result is a buffer consisting of the intermediate and basic species of oxalic acid. Use the HendersonHasselbalch equation with pK2 and 1.50 mmol HA− and 1.00 mmol A2− to calculate the pH. 26. 4.48; The addition of 25 mL of 0.125 0 M NaOH gives a buffer, with 3.125 mmol A− and 1.875 mmol HA. Use the Henderson-Hasselbalch equation to calculate the pH. 27. [NaHCO3] = 0.044 88 M; [Na2CO3] = 0.070 12 M; The first equivalence point corresponds to the conversion of the carbonate ion to the bicarbonate ion. Use the concentration and volume of HCl to determine the millimoles of HCl, which is equal to the millimoles of carbonate in the mixture. Divide by the volume of the mixture to determine the concentration. Subtract the volume of the first equivalence point from the second equivalence point to determine the volume of HCl that reacted with the bicarbonate in the solution. Determine the millimoles of HCl that corresponds to this volume. Subtract the millimoles of carbonate in the mixture to determine the millimoles of the bicarbonate that was in the original mixture. Divide by the volume of the mixture to determine the concentration.

28. 4.58% N; The millimoles of base added neutralizes the excess HCl. Subtract the millimoles of base from the total millimoles of HCl to determine the millimoles of HCl that reacted with the generate ammonium ion. For every mole of ammonium ion generated there is one mole of nitrogen in the sample. Therefore, the number of millimoles of HCl is equivalent to the number of millimoles of nitrogen in the sample. Convert millimoles to grams of nitrogen and divide by the sample mass. Multiply by 100 to convert to a percent.

29. 8.473; This is the equivalence point of the titration, so all of the weak acid has been converted to its conjugate base.

Calculate the concentration of the conjugate base and Kb for the conjugate base. Solve the weak-base equilibrium for [OH– ]. Then calculate the pOH and pH. 30. 1.745; This is before the first equivalence point of the titration. The result is a buffer consisting of the acidic and intermediate species of maleic acid. Use the Henderson-Hasselbalch equation with pK1 and 2.00 mmol HA− and 3.00 mmol H2A to calculate the pH. 31. 1.653; Determine the number of moles of NaOH added and the number of moles of HCl in the solution. The difference is the moles of HCl remaining in the solution. Determine the concentration of the HCl remaining in the solution using the total volume. Then calculate the pH.

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Chap 12_10e Indicate the answer choice that best completes the statement or answers the question.

1. A back titration is necessary if analyte I. precipitates in the absence of EDTA. II. reacts too slowly with EDTA. III. blocks the indicator. a. I b. II c. III d. I or II e. I, II, or III 2. There is not an indicator for the titration of Ag+ with EDTA. However, Ag+ can exchange Ni2+ from tetracyanonikelate(II). The titration of the liberated Ni2+ with EDTA can be used to determine the amount of Ag+ in the solution. This technique is called a. direct titration. b. indirect titration. c. displacement titration. d. back titration. e. masking.

3. Calculate pBa2+ after 35.00 mL of 0.100 M EDTA is added to 50.00 mL of 0.100 M Ba2+. The solution is buffered at pH 10. For the buffered pH of 10, = 0.30. Kf = 7.59 × 107 for BaY2−. a. 1.39 b. 1.75 c. 4.56 d. 4.82 e. 4.70 4. Calculate pBa2+ after 75.00 mL of 0.100 M EDTA is added to 50.00 mL of 0.100 M Ba2+. The solution is buffered at pH 10. For the buffered pH of 10, = 0.30. Kf = 7.59 × 107 for BaY2−. a. 7.06 b. 3.53 c. 4.33 d. 7.58 e. 4.59

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Chap 12_10e 5. The ___________ is the ability of multidentate ligands to form more stable metal complexes than those formed by similar monodentate ligands. a. ligand effect b. multidentate effect c. chelate effect d. Lewis effect e. isomer effect 6. Calculate pBa2+ after 50.00 mL of 0.100 M EDTA is added to 50.00 mL of 0.100 M Ba2+. The solution is buffered at pH 10. For the buffered pH of 10, = 0.30. Kf = 7.59 × 107 for BaY2−. a. 1.30 b. 7.88 c. 4.59 d. 7.36 e. 4.33 7. A(n) ____________ titration can be used to determine the concentration of anions that precipitate with certain metal ions by titrating the excess metal ion with EDTA. a. direct b. displacement c. indirect d. masking e. back 8. Which statement(s) regarding metal ion indicators in EDTA titrations is/are TRUE? I. Metal ion indicators are compounds that change color when they bind to a metal ion.

II. Useful indicators must bind the metal ion less strongly than EDTA does.

III. Metal ion indicators are also acid-base indicators. a. I only b. II only c. III only d. I and II e. I, II, and III

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Chap 12_10e 9. Which statement regarding auxiliary complexing agents is FALSE? a. Auxiliary complexing agents prevent the precipitation of metal hydroxide at high pH. b. Auxiliary complexing agents are chosen to bind the metal ion strongly enough to prevent metal hydroxide precipitation but weakly enough to be displaced by EDTA. c. is the effective formation constant at a fixed pH and a fixed concentration of the auxiliary complexing agent. d. A low concentration of the auxiliary complexing agent will obliterate the end point of the titration. e. = αMn+ αY4− Kf 10. Which is NOT an end point detection method used with EDTA titrations? a. metal ion indicators b. adsorption indicators c. glass (pH) electrode d. mercury electrode e. ion-selective electrode 11. A metal that does not freely dissociate from an indicator is said to ________ the indicator. a. hinder b. block c. restrict d. impair e. handicap 12. Which statement is INCORRECT for the conditional formation constant, Kf′? a. Kf′ is pH dependent. b. Kf′ = αY4–Kf c. Kf′ enables you to treat uncomplexed EDTA as if it were all in one form in equilibrium calculations. d. As pH decreases, the value of Kf′ increases. e. The value of Kf′ must be large to guarantee sharp end points during in EDTA titration.

13. A _____________ is a reagent that protects a component of the analyte from reaction with EDTA. a. hindrance agent b. displacement agent c. masking agent d. blocking agent e. reducing agent Copyright Macmillan Learning. Powered by Cognero.

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Chap 12_10e 14. A ligand that can bind to a metal ion through four ligand atoms is a ______________ ligand. a. monodentate b. bidentate c. tridentate d. tetradentate e. hexadentate 15. EDTA forms a ______ complex with all metal ions, regardless of the charge on the ion. a. 1:1 b. 1:2 c. 1:3 d. 1:4 e. 6:1 16. Ammonia is used as an auxiliary complexing agent for a Co2+ titration with EDTA. There are six β values for the complex between ammonia and Co2+. Which equation is used to calculate ? a.

b.

c.

d.

e.

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Chap 12_10e 17. For EDTA titrations, the titration reaction is Mn+ + EDTA ⇌ MYn−4. Which statements regarding EDTA titration curves are TRUE? I. At the equivalence point, there is exactly as much EDTA in solution as metal ion. [Mn+] is calculated from the dissociation of MYn−4 formed.

II. After the equivalence point, the concentration of free EDTA is determined from the excess EDTA, and virtually all metal ion is in the form MYn−4. [Mn+] is calculated from the dissociation of MYn−4.

III. Before the equivalence point, [Mn+] is determined from the unreacted Mn+ after the added EDTA has been consumed. Dissociation of MYn−4 is negligible.

IV. pM is plotted on the y-axis and volume of EDTA solution added is plotted on the x-axis. a. I, II, and III b. I, III, and IV c. III and IV d. I, II, III, and IV e. II, III, and IV 18. Calculate pBa2+ after 35.00 mL of 0.100 M EDTA is added to 50.00 mL of 0.100 M Ba2+ in the presence of 0.100 M nitrilotriacetate. The solution is buffered at pH 10. = 1.48 × 10−4. = 0.30 at pH 10.

Kf = 7.59 × 107 for BaY2−. a. 4.82 b. 5.58 c. 1.75 d. 4.55 e. 3.53 19. EDTA is a hexadentate ligand containing four carboxylic acid groups and two amines. Which statements regarding the acid-base properties of EDTA are TRUE? I. EDTA is a hexaprotic system, H6Y2+, when the amines are protonated.

II. Neutral EDTA, H4Y, in which the amines are not protonated, is tetraprotic.

III. Deprotonated EDTA, Y4−, is the only form of EDTA that binds to metal cations.

IV. The fraction of EDTA in any of its six forms is pH dependent. a. III and IV b. I and II c. I, III, and IV d. II, III, and IV e. I, II, and IV Copyright Macmillan Learning. Powered by Cognero.

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Chap 12_10e 20. Which statement is NOT true? a. A monodentate ligand binds to a metal atom through only one atom. b. A metal is multidentate if it can bind to more than one ligand. c. Metal ions are Lewis acids. d. Ligands are Lewis bases. e. Most transition metals bind six ligand atoms.

21. Calculate pBa2+ when 35.00 mL of 0.200 0 M EDTA is added to 50.00 mL of 0.100 0 M Ba2+ in the presence of 0.100 M nitrilotriacetate. = 1.48 × 10−4. = 0.30 at pH 10. Kf = 7.59 × 107 for BaY2−.

22. Calculate pMn2+ when 50.00 mL of 0.100 0 M Mn2+ is titrated with 25.00 mL of 0.200 0 M EDTA. The titration is buffered to pH 11, for which = 0.81. Kf = 7.76 × 1013 for MnY2−.

23. The reaction between Al3+ and EDTA is too slow for direct titration. To a 25.00-mL Al3+ solution of unknown concentration, 50.00 mL of 0.200 0 M EDTA is added and allowed to react for 30 min. The excess EDTA was titrated with 21.23 mL of 0.100 0 M Zn2+. What is the concentration of Al3+ in the sample?

24. A 25.00-mL solution containing an unknown amount of

was precipitated with 50.00-mL of 0.102 M

Ba2+ at pH 1. The BaSO4(s) was filtrated and washed. The excess Ba2+ in the filtrate required 21.03 mL 0.038 4 M EDTA to complete the complexometric titration. What was the concentration of in the original solution?

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Chap 12_10e 25. The formation constant of CoY2− is 1016.45. Calculate the concentration of free Co2+ in a solution of 0.10 M CoY2− buffered at pH 9.00. = 0.041 at pH 9.00.

26. Calculate pBa2+ when 25.00 mL of 0.200 0 M EDTA is added to 50.00 mL of 0.100 0 M Ba2+ in the presence of 0.100 M nitrilotriacetate. = 1.48 × 10−4. = 0.30 at pH 10. Kf = 7.59 × 107 for BaY2−.

27. Calculate pBa2+ when 15.00 mL of 0.200 0 M EDTA is added to 50.00 mL of 0.100 0 M Ba2+ in the presence of 0.100 M nitrilotriacetate. = 1.48 × 10−4. = 0.30 at pH 10. Kf = 7.59 × 107 for BaY2−.

28. Calculate pMg2+ when 25.00 mL of 0.100 M Mg2+ is titrated with 26.00 mL of 0.100 M EDTA. The titration is buffered to pH 10 for which = 0.30. Kf = 108.79 for MgY2−.

29. Calculate

for a 0.200 M EDTA solution buffered at pH 10.00. For EDTA, K1 = 1.0, K2 = 31.6, K3 =

100, K4 = 490, K5 = 1.35 × 106, and K6 = 2.34 × 1010.

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Chap 12_10e 30. Calculate pMn2+ when 50.00 mL of 0.100 0 M Mn2+ is titrated with 17.00 mL of 0.200 0 M EDTA. The titration is buffered to pH 11, for which = 0.81. Kf = 7.76 × 1013 for MnY2−.

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Chap 12_10e Answer Key 1. e 2. c 3. b 4. a 5. c 6. e 7. c 8. e 9. d 10. d 11. b 12. d 13. c 14. d 15. a 16. d 17. d 18. b 19. e 20. b 21. 3.13; Determine the millimoles of EDTA and the millimoles of Ba2+ to determine which one is in excess. There are more millimoles of EDTA than Ba2+ so essentially all Ba2+ is bound up in the complex. Since an auxiliary complexing agent is used, you must calculate the effective formation constant. Write the K expression, insert the concentrations of the complex and EDTA, and solve for the concentration of Ba2+. The assumption is made that is large enough that the change in complex concentration is negligible.

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Chap 12_10e 22. 7.49; The titration is at the equivalence point, so all of the Mn2+ reacts with all of the EDTA added to form the complex MnY2−. Divide the millimoles of Mn2+ by the total volume to get the concentration of the complex.

Calculate the conditional formation constant. Solve for the concentration of Mn2+ using the equilibrium expression.

Then calculate pMn2+.

23. 0.315 1 M Al3+; Calculate the total number of millimoles of EDTA added to the solution and the number of millimoles of Zn2+ required to react with the excess EDTA. The difference equals the millimoles of Al3+ in the 25.00-mL solution. Divide millimoles of Al3+ by 25.00 mL to determine the concentration Al3+.

24. 0.172 M

; Calculate the total number of millimoles of Ba2+ added to the solution and the number of millimoles

of EDTA required to react with the excess. The difference is equal to the number of millimoles of

in the

solution. Divide by the volume of the solution to determine the concentration.

25. 9.3 × 10−9 M Co2+; Determine the conditional formation constant. Solve the equilibrium for the formation of CoY2− for the concentration of Co2+. Assume the amount of CoY2− that dissociates is small relative to the initial concentration.

26. 2.352; The titration is at its equivalence point. Essentially all of the Ba2+ is part of the complex BaY2−. Calculate the concentration of the complex. Calculate the effective formation constant. Write the K expression, insert the concentration of the complex, and solve for the concentration of Ba2+. The assumption is made that is large enough that the change in complex concentration is negligible.

27. 5.342; There is excess Ba2+ in the solution at this point in the titration. Subtract the millimoles of EDTA added from the millimoles of Ba2+ to determine the excess. Divide by the total volume to determine the concentration of Ba2+.

Multiply by alpha for Ba2+ to determine the free concentration of Ba2+. Then calculate pBa2+.

28. 6.869; EDTA is in excess. Calculate the concentration of the excess EDTA and the concentration of the complex.

Then calculate the conditional formation constant. Use the equilibrium expression to calculate the concentration of free Mg2+. Then calculate pMg2+.

29. 0.299 0; The value of D is calculated using K1 through K6 and the hydronium concentration. The product of K1 through K6 is divided by the value of D to give the fraction of EDTA in the form Y4−.

30. 1.62; Find millimoles of EDTA and millimoles of Mn2+. Compare the number of millimoles to determine which is in excess (Mn2+). Divide the millimoles of excess Mn2+ by the total volume to get the concentration Mn2+ in solution.

Then calculate pMn2+.

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Chap 13_10e Indicate the answer choice that best completes the statement or answers the question.

1. Which is the correct effective equilibrium constant expression for the dissociation of −O2CCO2H in aqueous solution? ⇌

+ H+

a. b. c. d. e.

2. In addition to the equilibrium CaCO3(s) ⇌ Ca2+(aq) + (aq), additional equilibria can be written to give a more complete picture of the solubility of calcium carbonate. Which of the following equilibria is NOT valid? a. + H2O ⇌ + OH− Kb1 b.

+ H2O ⇌ H2CO3 + OH−

Kb2

c. Ca2+ + H2O ⇌ CaOH+ + H+

Ka

d. Ca2+ +

Kip

⇌ CaCO3

e. Ca2+ + H+ ⇌ CaH3+

Kip

3. Calculate the activity coefficient for phosphate in a solution with an ionic strength of 0.007 50 M.

a. 0.441 b. 0.140 c. 0.804 d. 0.913 e. 0.116

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Chap 13_10e 4. Calculate the pH for a buffer that is 0.10 M in NaO2CCO2Na and 0.50 M in NaO2CCO2H. K2 = 5.42 × 10−5,

= 0.333,

= 0.76,

a. 4.164 b. 4.567 c. 3.965 d. 4.323 e. 3.908 5. The mass balance equation for an aqueous solution of the sparingly soluble copper(II) carbonate is [Cu2+] = [ ] This mass balance equation does not take into account other equilibrium reactions that occur in solution. To determine a more accurate solubility for copper(II) carbonate, equilibria in addition to the dissolution reaction must be written for the copper(II) and carbonate ions in solution. Which mass balance equation is the only valid mass balance equation? a. [Cu2+] + [CuOH+] + [CuCO3] = [ ]+[ ] b. [Cu2+] + [CuCO3] = [

]+[

] + [H2CO3]

c. [Cu2+] + [CuOH+] = [C

]+[

] + [H2CO3] + [CuCO3]

d. [Cu2+] + [CuOH+] = [

]+[

] + [H2CO3]

e. [Cu2+] = [

]+[

]

6. The charge balance equation for an aqueous solution of the sparingly soluble cobalt sulfide is [H+] + 2[Co2+] = 2[S2−] + [OH−] This charge balance equation does not take into account other equilibrium reactions that occur in solution.

To determine a more accurate solubility for cobalt sulfide, equilibria in addition to the dissolution reaction must be written for the cobalt(II) and sulfide ions in solution. Which charge balance equation is the ONLY valid charge balance equation? a. [H+] + 2[Co2+] + [CoOH+] + 3[CoH3+] = 2[S2−] + [HS−] + [OH−] b. [H+] + 2[Co2+] + [CoOH+] = 2[S2−] + [HS−] + [OH−] + 3[SOH3−] c. [H+] + 2[Co2+] + [CoOH+] = 2[S2−] + [HS−] + [OH−] d. [H+] + 2[Co2+] + 3[CoH3+] = 2[S2−] + [HS−] + [OH−] + 3[SOH3−] e. [H+] + 2[Co2+] + [CoOH+] + 3[CoH3+] = 2[S2−] + [HS−] + [OH−] + 3[SOH3−]

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Chap 13_10e 7. What is the ionic strength of a buffer that is 0.15 M in NaHCO3 and 0.25 M in Na2CO3. a. 0.90 M b. 1.80 M c. 1.05 M d. 0.40 M e. 0.78 M 8. Which practice should NOT be used to maintain a nearly constant ionic strength in the titration of a diprotic acid (H2A) with NaOH? a. Add HCl to increase the degree of protonation of H2A. b. Add a fixed amount of a nonreactive salt, such as 0.10 M KCl, to the titration solution containing HCl and H2A. c. Maintain the concentrations of HCl and H2A at levels that are much less than the added 0.10 M KCl. d. The NaOH solution should be sufficiently concentrated so that the added volume is small relative to the initial volume of the acid mixture. e. These are all best practices for maintaining a nearly constant ionic strength in the titration.

9. The pH of a solution composed of NaO2CCO2H and NaCl needs to be determined. To get a more complete picture of the equilibria, the mass balance equation for the total oxalate is written. Which is the ONLY valid mass balance equation? a. Facid = [HO2CCO2H] + [HO2 ] + [– O 2 ] + [NaO2 ] + [NaO2CCO2Na] b. Facid = [HO2

] + [– O2

c. Facid = [HO2CCO2H] + [HO2 d. Facid = [NaO2

] + [NaO2 ] + [– O 2

] + [NaO2CCO2Na] ]

] + [NaO2CCO2Na]

e. Facid = [HO2CCO2H] + [HO2 [NaO2CCO2Na]

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] + [HO2CCO2Na] + [– O2

] + [NaO2

]+

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Chap 13_10e 10. Using the Davies equation, calculate the activities coefficient for Cu2+ in a solution with an ionic strength of 0.20 M.

a. 0.28 b. 0.31 c. 0.20 d. 3.2 e. 0.56 11. Consider an aqueous solution containing Fe(NO3)3, NaSCN, and HNO3 with the following chemistry.

Fe3+ + SCN− ⇌ Fe(SCN)2+ Fe3+ + 2SCN− ⇌ Fe3+ + H2O ⇌ FeOH2+ + H+ H2O ⇌ H+ + OH− Which mass balance equation is valid for this solution? a. [Fe3+] + [Fe(SCN)2+] + + [ ] = 3[ ] b. [Na+] = [SCN−] + [Fe(SCN)2+] + 2[

]

c. [H+] = [OH−] d. [Fe3+] + [H+] = [

]

e. [Fe3+] + [H+] = [

] + [OH−]

12. The mean fraction of protons,

, bound to H3A ranges from______________ and

(theoretical) =

______________. a. 0 to 3; b. 0 to 3; c. 0 to 3; d. 0 to 2; e. 0 to 2;

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Chap 13_10e 13. Many ions do not have tabulated activity coefficients, nor can the activity coefficient be estimated, as the ion size is now known. The Davies equation is used to estimate the activity coefficient under these conditions.

Which statement(s) regarding the Davies equation is/are FALSE? I. The calculated activity coefficient is independent of the size of the ion.

II. The charge of the ion, the ionic strength of the solution, and the size of the ion are required to calculate the activity coefficient.

III. The Davies equation can be used up to μ ≈ 0.5 M. IV. The calculated activity coefficient is no more accurate than guessing the ion size to determine the activity coefficient with the Debye-Hückel equation. a. II and IV b. II c. I d. III and IV e. III 14. The equilibrium constant that has had the activity coefficients at a given ionic strength incorporated into it is the a. conditional equilibrium constant, K'. b. temporary equilibrium constant, K'. c. combined equilibrium constant, K'. d. effective equilibrium constant, K'. e. activity equilibrium constant, K'.

15. Consider a saturated aqueous solution of AgCN(s) with the following chemistry.

AgCN(s) ⇌ Ag+ + CN− HCN ⇌ H+ + CN− Ag+ + H2O ⇌ AgOH(aq) + H+ H2O ⇌ H+ + OH− a. [Ag+] = [CN−] b. [H+] = [CN−] c. [Ag+] = [H+] d. [H+] = [OH−] e. [Ag+] + [H+] = [CN−] + [OH−]

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Chap 13_10e 16. Which of the fractional composition equations is INCORRECTLY written for the diprotic acid H2A and the monoprotic acid HA? a.

b.

c.

d.

e.

17. The pH of a solution composed of NaCH3CO2, NaO2CCO2H, and KCl needs to be determined. To get a more complete picture of the equilibria, the ion-pairing equilibria are included in the calculations. Which of the following is NOT a valid ion pairing for this solution? a. K+ +

b. Na+ +

⇌ KCH3CO2

⇌

c. Na+ + Cl− ⇌ NaCl

d. Na+ +

e.

⇌ NaCH3CO2

+ Cl− ⇌ CH3CO2Cl2−

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Chap 13_10e 18. Which is the correct effective solubility constant expression for the dissolution of PbCl2?

PbCl2(s) ⇌ Pb2+(aq) + 2Cl−(aq) a.

b.

c.

d.

e.

19. The equilibrium constants for a diprotic acid are determined from a Bjerrum plot by a. fitting the theoretical curve to the experimental curve by adjusting the pH to minimize the square of the residuals. b. determining the slope for the upper curve, which corresponds to K1, and the slope for the lower curve, which corresponds to K2. c. varying the pH to determine the range over which the slope of the theoretical curve is zero. The lower end of the range is K1 and the upper end of the range is K2. d. fitting the theoretical curve to the experimental curve by the method of least squares to find the values of K1 and K2 that minimize the square of the residuals. e. taking the first derivative of the experimental data to determine the minimums in the slope. Each minimum corresponds to an equilibrium constant.

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Chap 13_10e 20. A student is tasked with determining the pH of a solution prepared by dissolving 20.0 mmol sodium acetate, 10.0 mmol sodium oxalate (NaO2CCO2H), and 15.0 mmol potassium chloride in enough water to prepare 1.00 L of solution. Which of the following will simplify the algebra used to solve the problem?

I. Write fractional composition equations for each acid and base in the charge balance equation.

II. Substitute the fractional composition equations into the charge balance equation and enter the known values for [Na+], [K+], and [Cl−].

III. Assume the weakest acid will not contribute significantly to the pH of the solution.

IV. Use a spreadsheet to perform the calculations. a. IV only b. I, II, and III c. I, II, III, and IV d. II, III, and IV e. I, II, and IV 21. Calculate the concentration of Mn2+ in a saturated Mn(OH)2 solution at pH 8.00. Consider the following equilibria.

Mn(OH)2(s) Mn2+ + 2OH− Ks p = 10−12.8 Mn2+ + OH− H2O

MnOH+

H+ + OH−

K1 = 103.4 Kw = 10−14.00

22. A solution is 0.120 M sodium oxalate (NaO2CCO2H), 0.050 M chloroacetic acid (ClCH2CO2H), and 0.010 M KOH. Express the charge balance equation for the solution in terms of [H+].

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Chap 13_10e 23. Consider the following equilibria for a saturated solution of CaSO4.

CaSO4

Ca2+ +

Ks p = 2.40 × 10−5

Ca2+ + H2O

CaOH+ + H+

Ka = 2.00 × 10−13

Ca2+ +

CaSO4

Kip = 229

H+ +

K2 = 1.03 × 10−2

Express [Ca2+] in terms [H+] for the solubility of CaSO4. Assume all activity coefficients are unity.

24. Including activity coefficients, calculate the pH of a sulfurous acid buffer prepared from 0.010 M NaHSO3 and 0.030 M Na2SO3. K1 = 1.39 × 10−2 and K2 = 6.73 × 10−8 for sulfurous acid.

Ionic strength (m, M) Ion 0.001 0.005 0.01 0.05 0.1 Charge = ± 1 Activity coefficient (g) 0.964 0.928 0.902 0.82 0.775 Charge = ± 2

Activity coefficient (g) 0.867 0.742 0.665 0.455 0.37

25. A solution is made by dissolving 0.25 mol MgBr2 in 1.0 L water. The species in solution are Mg2+, Br−, MgBr+, and MgOH+. Write the mass balance and charge balance equations for the solution.

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Chap 13_10e 26. Including activity coefficients, calculate the pH of a trimethylamine buffer prepared from 0.010 M (CH3)3NHCl and 0.015 M (CH3)3NH. Ka = 1.5 × 10−10 for (CH3)3NH+. , .

27. A solution is 0.120 M in HNO2 and 0.050 M NaCl. Calculate the [H+] of the solution with and without activity coefficients. What is the percent error in the concentration when the activity coefficients are neglected? Ka = 7.1 × 10−4 for HNO2.

,

, and

.

28. Consider the equilibria involved in the CO2-carbonate system in aqueous solution.

CO2(g) + H2O H2CO3

H2CO3

H+ +

KH = 3.4 102 Ka1 = 4.5 10−7

+ OH−

Kb1 = 2.1 10−4

H2O

H+ + OH−

Kw = 1.0 10−14

CaCO3(s)

Ca2+ +

Ks p = 4.6 10−9

Determine the pH of this system at

0.000 39 atm.

+ H2O

29. Calculate the concentration of each fumaric acid (HO2CCHCHCO2H) species in a 0.175 M solution at pH = 5.000. K1 = 9.5 × 10−4 and K2 = 3.3 × 10−5.

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Chap 13_10e 30. A solution is 0.200 M in acetic acid and 0.100 M KCl. Calculate the concentrations of acetic acid and the acetate ion in solution at pH 4.500.

= 0.775,

= 1.00,

= 0.83, and Ka = 1.75 ×

10−5.

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Chap 13_10e Answer Key 1. d 2. e 3. a 4. a 5. d 6. c 7. a 8. e 9. e 10. b 11. b 12. b 13. b 14. d 15. e 16. d 17. e 18. c 19. d 20. e 21. 4.9 × 10−7 M

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Chap 13_10e

22.

or ; Solve the problem in

terms of the K values and [H+] or incorporate the K values and [H+] into the α values. The molar concentrations are inserted for [Na+] and [K+].

23.

This problem is time intensive. Use the mass balance equation to solve for the various

species in terms of [Ca2+]. The concentration ion pairing cancels out in the mass balance equation, leaving [Ca2+] + [CaOH+] = [ ]+[ ].

24. 7.33 (7.328); Determine the ionic strength from the concentrations of the two ionic salts, and then determine the activity coefficients for the acid and conjugate base. Insert K2, concentrations, and activity coefficients into the Henderson-Hasselbalch equation to solve for pH. It is assumed that the K values do not change with changes in ionic strength. If activity coefficients are neglected, the pH is 7.65 (7.649).

25. mass balance for magnesium: [Mg2+] + [MgBr+] + [MgOH+] = 0.25 M; mass balance for bromine: [Br−] + [MgBr+] = 0.50 M; charge balance: [Mg2+] + [MgBr+] + [MgOH+] + [H+] = [Br−] + [OH−] 26. 10.05; Insert Ka, concentrations, and activity coefficients into the Henderson-Hasselbalch equation to solve for pH. It is assumed that the K value does not change with changes in ionic strength.

27. 16%; Solve the weak acid equilibria for [H+] with and without activity coefficients. Determine the difference between the values and divide by the concentration calculated with activity coefficients. Then multiply by 100 to determine the percent error.

28. 8.2 29. [H2A] = 0.000 431 M, [HA−] = 0.041 0 M and [A2−] = 0.133 7 M; Calculate the fractional composition of the solution from the pH and K values. Multiply the formal concentration by the fraction of each species to determine the concentration of each species.

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Chap 13_10e 30. [CH3CO2H] = 0.1074 M, [ ] = 0.092 4 M; Calculate the fraction of HA and the fraction of A− and multiply each fraction by 0.200 M to determine the concentrations of acetate and acetic acid. The activity coefficients are incorporated into the Ka value to give an effective equilibrium constant, which is used in the fractional composition equations.

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Chap 14_10e Indicate the answer choice that best completes the statement or answers the question.

1. For the following reaction, the oxidizing agent is ________ and it _________ electrons, and the reducing agent is __________ and it __________ electrons.

8H+ + + 5Cu+ ⇌ Mn2+ + 5Cu2+ + 4H2O a.

; gains 5; Cu+; loses 1

b. H+; loses 1; Cu+; gains 1 c. H+; gains 1;

; loses 5

d. Cu+; loses 1;

; gains 1

e. Cu+; loses 1; H+; gains 1 2. Which statements are TRUE when an electrochemical cell is used to determine the concentration of an analyte?

I. The concentrations of all species in the half-cells must be known—except for the analyte of interest.

II. The analyte concentration is calculated from the cell potential once the reaction reaches equilibrium.

III. The concentrations of all species in the half-cells remain largely unchanged because a negligible amount of current flows through the potentiometer.

IV. Half-cells other than the S.H.E. can be used to determine pH. a. II, III, and IV b. I, III, and IV c. I and II d. III and IV e. I, II, and IV 3. Eo' is defined as a. the formal half-cell potential for compounds that are neither acids nor bases. b. the formal half-cell potential at 20°C. c. the formal half-cell potential at pH 7. d. the inverse of Eo . e. None of these answers is the definition of Eo'.

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Chap 14_10e 4. The following half-cells are connected via a salt bridge and potentiometer to form a galvanic cell. What is the voltage read by the potentiometer when AAg+ = 2.00 and AV2+ = 1.50?

Ag+(aq) + e− → Ag(s) V2+(aq) + 2e− → V(s)

E° = 0.799 3 V E° = −1.125 V

a. 1.912 V b. −0.326 V c. −0.313 V d. 1.937 V e. 1.924 V 5. Which relationship is NOT correctly defined? a. Work = E · q b. ΔG = −n · F · E c. d. P = E ∙ I e. q = n ∙ N ∙ F 6. The formal potential is the reduction potential that applies under a specified set of conditions, including I. pH. II. ionic strength.

III. concentration of complexing agents. a. I b. II c. II and III d. I and II e. I, II, and III 7. For the half-reaction Ag+ + e− → Ag(s), when the concentration of Ag+ is increased, the reduction potential a. becomes more positive. b. becomes more negative. c. remains constant. d. increases or decreases depending on the voltage of the other half-reaction. e. increases or decreases depending on the temperature.

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Chap 14_10e 8. Find the standard cell potential for Cu(s) + 2Ag+ 2Ag(s) + Cu2+.

Cu2+ + 2e− ⇌ Cu(s) E° = 0.339 V + − Ag + e ⇌ Ag(s) E° = 0.799 V a. −0.460 V b. 0.460 V c. 1.259 V d. 0.630 V e. −1.259 V 9. Consider the following line diagram for an electrochemical cell. If the measured potential is 1.00 V, what is the pH of the dichromate cell?

Fe | Fe2+(aq, 1.00 M) || (aq, 1.00 M), Cr2+(aq, 1.00 M), H+(aq, x M) | Pt + 14H+ + 6e− ⇌ 2Cr3+ + 7H2O Fe2+ ⇌ + 2e− Fe(s) a. 1.931 b. 5.795 c. 12.162 d. 6.081 e. 4.054

E° = 1.36 V E° = −0.44 V

10. A 2.35-g sample of Zn (FM 65.38) was oxidized according to the following reaction.

Zn(s) + 2H+(aq) ⇌ Zn2+(aq) + H2(g) How many coulombs of charge are transferred from Zn to H+ when the sample is completely oxidized? a. 6.94 × 103 C b. 3.47 × 103 C c. 1.39 × 104 C d. 2.78 × 104 C e. 1.74 × 103 C

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Chap 14_10e 11. Calculate the equilibrium constant for the reaction between Sn(s) and Zn2+(aq).

Sn(s) + Zn2+(aq) ⇌ Sn2+(aq) + Zn(s) Sn2+ ⇌ + 2e− Sn(s) E° = −0.141 V Zn2+ ⇌ + 2e− Zn(s) E° = −0.762 V a. 9.86 × 1020 b. 1.201 × 10−21 c. 3.37 × 1030 d. 3.18 × 10−11 e. 1.84 × 1015 12. A galvanic cell with a large equilibrium constant has a _________ cell potential. a. negative b. constant c. zero voltage d. positive e. variable 13. A galvanic cell is at equilibrium when a. Ecell is a constant value. b. the cell resistance is very low. c. Ecell = 0 V. d. Ecell fluctuates between

.

e. the cell resistance is infinitely large.

14. Which statement regarding galvanic cells is FALSE? a. Galvanic cells are spontaneous. b. Oxidation occurs at the anode, and reduction occurs at the cathode. c. Electrons move toward the more negative electrical potential. d. Galvanic cells are composed of two half-cells connected by a salt bridge. e. The salt bridge maintains electroneutrality throughout the cell.

15. What is the potential change (in volts) when 1.24 mmol of electrons do 45.5 J of work? a. 0.435 V b. 0.213 V c. 0.013 V d. 0.564 V e. 0.380 V

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Chap 14_10e 16. Which species is the strongest oxidizing agent?

O3(g) + 2H+ + 2e− ⇌ O2(g) + H2O E° = 2.075 V Ag+ + e− ⇌ Ag(s) Li+ + e− ⇌ Li(s) 2H+ + 2e− ⇌ H2(g) + 8H+ + 5e− ⇌ Mn2+ + 4H2O

E° = 0.799 V E° = −3.040 V E° = 0.00 V E° = 1.507 V

a. O3 b. Ag+ c. Li+ d. H+ e.

17. Which statement regarding the following line notation is FALSE? Fe(s) | Fe2+ (aq, A = 1) || Cu2+ (aq, A = 1) | Cu(s) a. Fe is the anode. b. | represents a phase boundary. c. There is no salt bridge in this cell. d. 1 M Cu2+ is the electrolyte concentration and species in the cathode half-cell. e. 1 M Fe2+ is the electrolyte concentration and species in the anode half-cell.

18. Which statement(s) regarding the measurement of standard reduction potentials is/are FALSE? I. Standard conditions for measuring reduction potentials are 25oC and A = 1 for all aqueous and gaseous species.

II. Reduction potentials are measured relative to the standard hydrogen electrode (S.H.E.).

III. Standard reduction potentials are measured while the half-cell is connected to the negative terminal of the potentiometer.

IV. The E° for the S.H.E. is defined as 0 V. a. I and IV b. I and III c. I d. III e. II and IV

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Chap 14_10e 19. Which statement(s) is/are NOT correct when the silver and vanadium half-cells are connected via a salt bridge and a potentiometer to form a galvanic cell?

Ag+(aq) + e− → Ag(s) E° = 0.799 3 V V2+(aq) + 2e− → V(s) E° = −1.125 V I. Ag+ is reduced.

II. V is oxidized.

III. = 1.924 V IV. V2+ is reduced.

V. Ag is oxidized. a. I and II b. III, IV, and V c. I, II, and III d. III only e. IV and V 20. Calculate Eo' for the following half-reaction.

O2 + 4H+ + 4e− ⇌ 2 H2O Eo = 1.229 V a. 1.643 V b. 0.815 V c. 0.307 V d. 1.229 V e. 0.711 V 21. Find ΔGo and K for the reaction + e−

E° = 0.017 V E° = 0.356 V

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Chap 14_10e 22. Calculate the voltage that would be measured for the following reaction if ΔG is −45.0 kJ/mol I2.

2Fe3+ + 2I− → 2Fe2+ + I2

23. A bromide probe is constructed from a saturated silver bromide half-cell and a Ni2+/Ni half-cell. If the measured potential is 0.952 V and [Ni2+] = 1.00 M, what is the bromide concentration?

AgBr(s) + e−

Ni2+ + 2e−

Ag(s) + Br− E° = 0.071 V

E° = − 0.714 V

Ni(s)

24. Calculate the standard potential when the two half-cells are connected via a salt bridge and potentiometer to form a galvanic cell. Which species is the reducing agent and which species is the oxidizing agent?

E° = 1.589 V Mg(OH)2(s) + 2e−

Mg(s) + 2OH−

E° = −2.690 V

25. Use the Latimer diagram to calculate E° for Cr(III) → Cr(0).

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Chap 14_10e 26. What current is required to reduce Al3+ to Al(s) at a rate of 1.50 mmol/h?

27. Calculate Eo and K for a galvanic cell composed of a bromate/bromide half-cell and a chlorate/chlorine dioxide half-cell.

ClO2 + H2O E° = 1.130 V Br− + 6OH−

E° = 0.613 V

28. For the following cell diagram, write and balance the reaction that occurs in the cell.

Pt | HOCl(aq, 1 M), HClO2(aq, 1 M), H+ (aq, 1 M) || (aq, 1 M), H+(aq, 1 M) | MnO2(s) | Pt

29. Compute Eo ´ for the half-reaction

HNO2 + H2O. E° = 0.940 V and Ka = 7.2

10–4.

30. Calculate the potential for the following half-reaction when [

] = 0.100 M, [

] = 0.500 M, and

the p3 H = 10.50.

E° = −0.566 V

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Chap 14_10e Answer Key 1. a 2. b 3. c 4. d 5. a 6. e 7. a 8. b 9. b 10. a 11. b 12. d 13. c 14. c 15. e 16. a 17. c 18. d 19. e 20. b 21. ΔGo = 32.7 kJ and K = 1.9 × 10−6; Calculate Eo . Then use the equations ΔGo = nFEo and

.

22. 0.233 V; Use the equation that relates ΔG and E. There are two electrons transferred per atom of I2 in the reaction.

23. [Br−] = 1.50 × 10−3 M; Determine E°. Insert E° and [Ni2+] into the Nernst equation to solve for [Br−].

24. 4.279 V;

is the oxidizing agent and Mg(s) is the reducing agent.

25. −0.740 V; The sum of the ΔG expressions for the two half-reactions for which the potentials are known is equal to the ΔG expression for Cr(III) → Cr(0).

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Chap 14_10e 26. 121 mA; Convert the rate from millimoles of Al per hour to millimoles of electrons per hour. Then use the Faraday constant to convert the rate to coulombs per hour. Finally, convert the rate to coulombs per second, which is equivalent to amperes, which is the unit for current.

27. E° = 0.517 V and K = 2.72 × 1052; Subtract the standard reduction potential of the bromate/bromide half-reaction from the standard reduction potential of the chlorate/chlorine dioxide half-reaction. Insert E° and n = 6 into .

28.

2MnO2 + 3HClO2 + H2O

29. 0.433 V 30. −0.281 V; Use the Nernst equation, plugging in the concentration values and the moles of electrons. Calculate the OH− from the pH.

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Chap 15_10e Indicate the answer choice that best completes the statement or answers the question.

1. Which is NOT a class of ion-selective electrode? a. glass membranes b. solid-state electrodes c. liquid-based electrodes d. compound electrodes e. biochemical electrodes 2. Solid-state chemical sensors are manufactured using semiconductors such as Si and Ge. Doping enhances the conductivity of Si. Which is NOT true for doping of silicon? a. Doping replaces some of the silicon atoms with atoms of different elements. b. Doping silicon with phosphorous enhances the conductivity of silicon by providing an extra conduction electron that is free to move through the crystal. c. Doping silicon with aluminum enhances the conductivity of silicon by creating a vacancy in the crystal called a hole, which behaves as a negative charge carrier. d. n-Type silicon has excess conduction electrons. e. p-Type silicon has excess holes.

3. When the concentration of H+ ions is very low and the concentration of Na+ ions is high, the apparent pH measured by a glass pH electrode is a. lower than the true pH of the solution. b. higher than the true pH of the solution. c. the same as the true pH of the solution. d. equal to twice the concentration of the Na+ ions. e. equal to half the concentration of the Na+ ions.

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Chap 15_10e 4. Which statement(s) is/are NOT true for an ion-selective electrode with an organic polymer membrane specific to aqueous ions?

I. Two reference electrodes on either side of the organic polymer membrane of the ion-selective electrode measure the potential difference across the membrane.

II. The organic polymer membrane of the ion-selective electrode separates the inner solution from the outer analyte solution.

III. The organic polymer membrane of the ion-selective electrode is hydrophobic and is impregnated with a viscous organic solution containing an ion exchanger and sometimes a ligand specific for the analyte ion.

IV. The potential difference across the organic polymer membrane of the ion-selective electrode occurs when the analyte ions embedded in the membrane diffuse into the inner aqueous solution, creating a junction potential.

V. The potential difference across the organic polymer membrane of the ion-selective electrode is dependent only on the activity of analyte ion in the outer solution. a. IV b. I and V c. II and III d. IV and V e. V 5. The ______________ is the voltage difference when two dissimilar electrolyte solutions are in contact. a. ion potential b. electrolyte potential c. cation potential d. junction potential e. anion potential 6. An electrode with a fixed potential is the a. cathode. b. anode. c. indicator electrode. d. reference electrode. e. working electrode.

7. To recover a glass pH electrode that is responding sluggishly, you should soak it for a few minutes in a. water. b. 20 wt% aqueous ammonium bifluoride. c. 10 wt% aqueous ammonia. d. 20 wt% aqueous lithium perchlorate. e. 10 wt% aqueous sodium chloride.

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Chap 15_10e 8. The potential of a silver electrode is measured relative to a saturated silver-silver chloride reference electrode for the titration of 100.0 mL of 0.100 M Cl- with 0.100 M Ag+. What is the potential after 75.00 mL of titrant is added? E°(Ag+/Ag(s)) = 0.799 V, E(saturated KCl) = 0.197 V for the Ag|AgCl electrode, and Ks p(AgCl) = 1.8 × 10−10. a. 0.493 V b. 1.070 V c. 0.521 V d. 1.267 V e. 0.135 V 9. A Pb2+ ion-selective electrode is moved from a 6.41 × 10−5 M Pb(NO3)2 solution to a 4.84 × 10−3 M Pb(NO3)2 solution at 25°C. By how many millivolts will the potential of the ion-selective electrode change when it is moved from the first solution to the second solution? a. 55.55 mV b. −55.55 mV c. 111.1 mV d. −111.1 mV e. 59.16 mV 10. Which is NOT a potential error associated with a pH measurement that is incorrectly defined? a. Junction potential drift—AgCl or Ag precipitates from solution in the porous plug due to dilution of the KCl or the presence of reducing agents. b. Sodium error—The electrode responds to Na+ instead of H+ when [Na+] is low and [H+] is high. c. Acid error—In strong acid, the measured pH is higher than the actual pH because the glass is saturated with H+ and cannot be further protonated. d. Temperature—The pH meter is calibrated at a different temperature than that at which the measurement is made. e. Equilibration time—The electrode is not given sufficient time to equilibrate with the solution.

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Chap 15_10e 11. An Na+ ion-selective electrode increased in voltage by 59.16 mV when the sodium ion activity increased by a factor of 10. An S2− ion-selective electrode measured a decrease of 29.58 mV when the sulfide activity increased by a factor of 10. Why? I. For cations, the measured potential increases by a factor of n × 59.16 mV for every 10-fold change in the activity of the cation, where n is the charge on the cation.

II. For anions, the measured potential decreases by a factor of n × 59.16 mV for every 10-fold change in the activity of the anion, where n is the charge on the anion.

III. For cations, the measured potential increases by a factor of 59.16 mV/n for every 10-fold change in the activity of the cation, where n is the charge on the cation.

IV. For anions, the measured potential decreases by a factor of 59.16 mV/n for every 10-fold change in the activity of the anion, where n is the charge on the anion. a. II and III b. III and IV c. I and II d. I and IV e. None of the statements explains the behavior.

12. When using ion-selective electrodes, to compensate for a complex or unknown matrix, the _____________ method can be used to determine the analyte concentration. a. calibration curve b. standard addition c. standardization d. least-squares analysis e. dilution 13. The main advantage to using a saturated KCl solution in reference electrodes is a. the low cost of potassium chloride. b. the ease of reaction with analyte. c. that the chloride concentration does not change as liquid evaporates from the cell. d. the ease of cell construction. e. the ease at which chloride passes through membranes.

14. The ________________ gives the relative response of an ion-selective electrode to different species with the same charge. a. counter ion coefficient b. isoionic coefficient c. ion coefficient d. selectivity coefficient e. transient coefficient Copyright Macmillan Learning. Powered by Cognero.

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Chap 15_10e 15. An electrode has a potential of –0.335 V with respect to a saturated calomel electrode (S.C.E.). What is the potential with respect to a saturated silver–silver chloride electrode (Ag|AgCl)? E(saturated KCl) = 0.197 V for the Ag|AgCl electrode and E(saturated KCl) = 0.241 V for the S.C.E. a. −0.291 V b. 0.291 V c. 0.532 V d. −0.532 V e. 0.044 V 16. Solid-state ion-selective electrodes that respond to anions create a potential difference across the solidstate membrane by ___________________, which allows anions to migrate through the crystal. a. doping the membrane crystal to create anion vacancies b. drilling small channels through the crystal c. reacting the membrane crystal with a strong acid d. polishing the surface of the membrane crystal e. etching the surface of the membrane crystal with small grooves 17. Which statements are TRUE for indicator electrodes?

I. Indicator electrodes are made from relatively inert metals.

II. Indicator electrodes work best when the electrode surface is large and clean.

III. Platinum is a common indicator electrode.

IV. Platinum electrodes are more inert than gold electrodes. a. I, III, and IV b. II and III c. I, II, and IV d. I, II, and III e. I, II, III, and IV

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Chap 15_10e 18. A 0.100 M KCl solution is placed in contact with a 0.100 M LiCl solution. Which statements are TRUE for the junction between the two solutions? The mobilities of the ions are K+ = 7.62 × 10−8 m2/(s∙V), Li+ = 4.01 × 10−8 m2/(s∙V), and Cl−= 7.91 × 10−8 m2/(s∙V).

I. The KCl side of the junction will be positive as K+ diffuses more quickly across the junction than Cl −.

II. The LiCl side of the junction will be negative as Li+ diffuses more quickly across the junction than K+.

III. The KCl side of the junction will be positive as K+ diffuses more slowly across the junction than Li+.

IV. The LiCl side of the junction will be negative as Li+ diffuses more quickly across the junction than Cl −. a. I and III b. II and IV c. I and II d. I and IV e. II and III 19. Which method is NOT a means to lower the detection limit of a liquid-based ion-selective electrode? a. reduce the concentration of the primary (analyte) ion in the inner filling solution with a metal ion buffer b. lower the mobility of the primary ion in the ion-selective membrane so that the primary ion cannot leak from the inner solution c. replace the inner filling solution with an electrically conductive polymer d. use a nonporous frit in place of the hydrophobic polymer membrane e. All of the methods can be used to lower the detection limit.

20. An electrode has a potential of 1.201 V with respect to a saturated silver-silver chloride (Ag|AgCl) electrode. What is the potential with respect to a saturated calomel electrode (S.C.E.)? E(saturated KCl) = 0.197 V for the Ag|AgCl electrode and E(saturated KCl) = 0.241 V for the S.C.E. a. 1.442 V b. 1.245 V c. 1.004 V d. 1.398 V e. 1.157 V

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Chap 15_10e 21. The selectivity coefficient for a fluoride ion-selective electrode is

= 0.15. What will the change in

the electrode potential be if the pH of a 1.4 × 10−5 M F− solution is raised from pH 5.50 to pH 11.00?

22. A cell is prepared by dipping a wire of metal X(s) and a saturated calomel electrode (E(saturated KCl) = 0.241 V) into a 0.250 M X2+ solution. The X(s) wire is attached to the positive terminal of the potentiometer, and the saturated calomel electrode is attached to the negative terminal of the potentiometer.

The cell potential was found to be 0.137 V. Calculate E° for the half-reaction X2+(aq) + 2e− → X(s).

23. A Cu2+ ion-selective electrode was calibrated using metal ion buffers with a constant ionic strength. The response of the electrode to these buffer solutions is shown in the table. Additionally, the response of the electrode to a solution containing Cu2+ and Mg2+ is shown in the table. Calculate the selectivity coefficient for Mg2+, , for this Cu2+ ion-selective electrode.

[Cu2+] (M) 2.14 × 10−6 5.59 × 10−4 2.14 × 10−6

[Mg2+] (M) 0 0 1.17 × 10−2

Response (mV) −66.82 1.40 −34.88

24. When an ion-selective electrode for X+ was immersed in 0.032 0 M XCl, the measured potential was 0.042 0 V. What is the concentration of X+ in a solution with a measured potential of 0.057 0 V? Assume that the electrode follows the Nernst equation, the temperature is 25°C, and the activity coefficient of X+ is 1.00.

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Chap 15_10e 25. A pH probe is calibrated against pH 4 and pH 7 buffers. For the pH 4 buffer, E = 200 mV and for the pH 7 buffer, E = 23 mV. The potential for a solution of unknown pH is 157 mV. Calculate the pH of the solution.

26. Calculate the pH change of a glass electrode when the activity of H+ in the solution increases by 5.0%.

27. The measured potential of a half-cell is 0.584 V when measured against the saturated calomel electrode (S.C.E.). What is the potential if the measurement is made using a silver-silver chloride (Ag|AgCl) reference electrode? E(saturated KCl) = 0.197 V for the Ag|AgCl electrode and E(saturated KCl) = 0.241 V for the S.C.E.

28. The potential of a silver wire electrode is measured relative to a saturated silver-silver chloride reference electrode for the titration of 25.00 mL of a 0.150 M Cl− solution with 0.100 M Ag+. Calculate the potential after 50.00 mL of 0.100 M Ag+ is added. E°(Ag+/Ag(s)) = 0.799 V, E(saturated KCl) = 0.197 V for the Ag|AgCl electrode, and Ks p(AgCl) = 1.8 × 10−10.

29. The potential of a silver wire electrode is measured relative to a saturated silver-silver chloride reference electrode for the titration of 50.00 mL of a 0.150 M Cl− solution with 0.100 M Ag+. Calculate the potential after 50.00 mL of 0.100 M Ag+ is added. E°(Ag+/Ag(s)) = 0.799 V, E(saturated KCl) = 0.197 V for the Ag|AgCl electrode, and Ks p(AgCl) = 1.8 × 10−10.

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Chap 15_10e 30. Determine which side of the liquid junction 0.1 M NaNO3 | 0.1 M KI will be positive. The mobilities of the ions are K+ = 7.62 × 10−8 m2/(s∙V), Na+ = 5.19 × 10−8 m2/(s∙V), = 7.40 × 10−8 m2/(s∙V), and Cl− = 7.91 × 10−8 m2/(s∙V).

31. A calomel electrode has the potentials E° = 0.268 V and E(saturated KCl) = 0.241 V. Use these potentials to calculate the activity of Cl− in saturated KCl. Then calculate E for a silver-silver chloride electrode saturated with KCl. E° = 0.222 V for the silver-silver chloride electrode.

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Chap 15_10e Answer Key 1. e 2. c 3. a 4. a 5. d 6. d 7. b 8. e 9. a 10. b 11. b 12. b 13. c 14. d 15. a 16. a 17. d 18. e 19. d 20. e 21. The potential decreases by 63.23 mV, or DE = −63.23 mV or −0.063 23 V. Calculate the potential at pH 5.5, where [OH−] is negligible and the selectivity coefficient can be ignored. Then calculate the potential at pH 11.0, where [OH−] is not negligible, and use the selectivity coefficient to correct for OH−. The potential difference is EpH 11.0 – EpH 5.50.

22. 0.360 V 23. 2.20 × 10−3 24. 0.057 4 M Copyright Macmillan Learning. Powered by Cognero.

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Chap 15_10e 25. pH = 4.73; Calculate ΔE/ΔpH for the buffers and then solve for pHunknown when the potential Eunknown is measured using the relationship ΔE/ΔpH = (Eunknown – ES1)/(pHunknown – pHS1).

26. ΔpH = 0.021; E = constant + b(0.059 16) log 1.00; E = constant + b(0.059 16) log 1.05; DE = b(0.059 16) log (1.00/1.05) = 1.3 mV; 1.3 mV = 1.3 × 10−3 V = b(0.059 16) V/pH unit; ΔpH = 0.021 27. 0.628 V ; The potential is positive, so E for the half-cell is more positive than E for the S.C.E. 0.584 V = x - 0.241 V, x = 0.825 V. Next, calculate the potential relative to Ag|AgCl by 0.825 V - 0.197 V = 0.625 V. 28. 0.497 V; Calculate millimoles of Ag+ and Cl− after the addition and determine which species is in excess (Ag+).

Calculate the concentration of Ag+ using the millimoles of excess Ag+ and the total volume. Plug into the Nernst equation (E = Eo + (0.059 16/n)log[Ag+] − Eref). Solve for E. 29. 0.120 V; Calculate millimoles of Ag+ and Cl− after the addition and determine which species is in excess (Cl−).

Calculate the concentration of Cl− using the millimoles of excess Cl− and the total volume. Use the Ks p to determine [Ag+] and plug into the Nernst equation (E = Eo + (0.059 16/n)log[Ag+] − Eref). Solve for E. 30. The left side (0.1 M NaNO3) of the junction will be positive.

31. 0.195 V

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Chap 16_10e Indicate the answer choice that best completes the statement or answers the question.

1. What isthe potential of the indicator electrode (versus the reference electrode potential) at the equivalence point of the titration of Sn2+ with Fe3+ to give Sn4+ and Fe2+ in 1 M HCl? E°(Fe3+/Fe2+) = 0.732 V; E° (Sn4+/Sn2+) = 0.139 V a.

b.

c.

d.

e.

2. Consider the following redox reaction. → A 0.588 2-g sample of Fe(NH4)2(SO4)2 ∙ 6H2O (FM 392.14) is dissolved in 400. mL of 1.0 M H2SO4 and titrated with 0.020 0 M KMnO4. Calculate the potential at the equivalence point versus a saturated calomel reference electrode (E(saturated KCl) = 0.241 V).

Fe3+ + e− ⇌ Fe2+ Eo = 0.68 V → Eo = 1.507 V a. E = 1.128 V b. E = 0.853 V c. E = 0.577 V d. E = 1.369 V e. E = 1.094 V

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Chap 16_10e 3. What is the potential of the indicator electrode (versus the reference electrode potential) when the volume of added titrant is equal to one-half the volume of titrant required to reach the equivalence point? a. the standard reduction potential for the titrant half-reaction b. half the standard reduction potential for the titrant half-reaction c. half the sum of the standard reduction potentials for the analyte and titrant half-reactions d. the standard reduction potential for the analyte half-reaction e. half the standard reduction potential for the analyte half-reaction 4. Sodium thiosulfate is used almost exclusively to titrate with sodium thiosulfate? a. In neutral or acidic solution, b. Each mole of

. Which statement is FALSE for the titration of

oxidizes thiosulfate to tetrathionate.

that reacts with thiosulfate is equivalent to two moles of I2.

c. In basic solution,

disproportionates to I− and HOI, which can oxidize

d. Sodium thiosulfate is standardized with freshly made

to

.

solution prepared from KIO3 and KI.

e. Dissolved CO2 and metal ion–catalyzed atmospheric oxidation of thiosulfate decreases the concentration of a prepared thiosulfate solution.

5. Which of the statements are TRUE regarding the use of indicators with dichromate?

I. In acidic solution, dichromate is orange and changes color to green when reduced. The color change is abrupt enough to serve as an indicator.

II. The equivalence point for a redox titration involving dichromate can be monitored using Pt and calomel electrodes.

III. Dichromate end points are determined using the distinct color changes of diphenylamine sulfonic acid or diphenylbenzidine sulfonic acid indicators.

IV. In basic solution, dichromate is orange and changes color to yellow when reduced to the chromate anion. The color change is abrupt enough to serve as an indicator. a. II and III b. I and IV c. I and II d. I and III e. II and IV 6. ________________ is the titration of iodine produced by a chemical reaction. a. Iodigravimetric titration b. Iodimetry c. Iodogravimetric titration d. Iodometry e. Iodotitration

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Chap 16_10e 7. Consider the titration of Ce4+ with Cu+ to form Ce3+ and Cu2+. At what volume of titrant added (V) is the potential of the indicator electrode equal to the standard reduction potential of the titrant halfreaction? Ve represents the equivalence point volume. a.

b.

c. V = Ve d. V = 2Ve e.

8. Consider the titration of 50.00 mL of 0.100 M Co3+ with 0.130 M mL of 0.130 M

. What is the potential after 35.00

has been added? The potential is measured against the Ag|AgCl reference electrode

(E(saturated KCl) = 0.197 V).

Co3+ + e− ⇌ Co2+ + e− ⇌

Eo = 1.92 V Eo = 0.56 V

a. 0.42 V b. 0.62 V c. 1.66 V d. 1.86 V e. 1.24 V

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Chap 16_10e 9. Potassium permanganate is a strong oxidant commonly used in redox titrations. Which of the following are TRUE for the titration chemistry of permanganate?

I. In strongly acidic solution, the permanganate half-reaction is

.

II. In neutral or alkaline solution, the permanganate half-reaction is .

III. In strongly alkaline solution, the permanganate half-reaction is

.

a. None of the half-reactions is true for permanganate under the given conditions. b. II and III c. I and III d. I and II e. I, II, and III 10. A Ce4+ standard solution is prepared with primary-standard-grade ammonium hexanitratocerate(IV), (NH4)2Ce(NO3)6. Which preparation of the solution will result in the concentration of Ce4+ remaining relatively constant over time? a. a Ce4+ solution prepared using distilled water b. a Ce4+ solution prepared using hot HCl solution c. a Ce4+ solution prepared using H2SO4 solution d. a Ce4+ solution prepared using HClO4 solution e. a Ce4+ solution prepared using HNO3 solution 11. Which statements are TRUE for redox indicators?

I. The color change occurs when the redox indicator goes from its oxidized state to its reduced state.

II. The color of the oxidized state of the redox indicator is observable when the ratio of the reduced state to the oxidized state is 1:10 or less.

III. The potential range over which a redox indicator changes is E = Eo ± 0.059 16/n.

IV. The color of the reduced state of the redox indicator is observable when the ratio of the reduced state to the oxidized state is 1:10 or less. a. I, II, III, and IV b. III and IV c. I, III, and IV d. I, II, and IV e. I, II, and III

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Chap 16_10e 12. Which statements are TRUE for the titration of Fe2+ with Ce4+ to give Fe3+ and Ce3+?

I. Before the equivalence point, the indicator electrode potential is calculated using the Fe3+/Fe2+ halfreaction because both [Fe3+] and [Fe2+] are known.

II. At the equivalence point, the indicator electrode potential is the average of Eo for the Fe3+/Fe2+ halfreaction and the Ce4+/Ce3+ half-reaction.

III. After the equivalence point, the indicator electrode potential is calculated using the Ce4+/Ce3+ halfreaction because both [Ce4+] and [Ce3+] are known. a. I and II b. II and III c. I, II, and III d. I and III e. None of the statements are true.

13. What is the potential range (versus S.C.E.; E(saturated KCl) = 0.241 V) over which the redox indicator phenosafranine (E° = 0.280 V versus S.H.E.) will change color? a. 0.020 to 0.098 V b. 0.039 to 0.521 V c. 0.221 to 0.339 V d. 0.079 to 0.157 V e. 0.009 5 to 0.068 5 V 14. Why is potassium permanganate NOT a primary standard? a. Solid KMnO4 contains traces of MnO2. b. Permanganate reacts with organic material in distilled water. c. Permanganate is unstable in aqueous solution, producing MnO2, O2, and OH−. d. Heat and light increase the rate of reaction between permanganate and water. e. All of the above are reasons potassium permanganate is not a primary standard.

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Chap 16_10e 15. In addition to redox indicators and the starch-iodine complex, the equivalence point for a redox titration can be determined using a Gran plot. Which statement is NOT true for a Gran plot? a. A Gran plot gives the best result when it uses titration data from well before the equivalence point. b. A Gran plot is constructed by plotting the volume of titrant added (V) on the x-axis and V·10−nE/0.059 16 on the y-axis.

c. A plot of V versus V·10−nE/0.059 16 is a straight line with the x-intercept equal to the equivalence point volume of the titration. d. A constant ionic strength over the course of the titration minimizes the error in the equivalence point. e. A Gran plot is more error prone than a first derivative or second derivative plot.

16. Sometimes the oxidation state of the analyte needs to be adjusted before it is titrated. The analyte can be preoxidized with an oxidizing agent to quantitatively oxidize the analyte. The excess oxidizing agent is then eliminated to prevent interference with the titration. Which oxidizing agent is paired with an INCORRECT elimination chemistry? a. Excess peroxydisulfate is eliminated by boiling, which converts the peroxydisulfate to sulfate. b. Excess silver(I,III) oxide in mineral acid is eliminated by boiling, which converts silver(III) to silver(I). c. Excess solid sodium bismuthate is removed by filtration. d. Excess hydrogen peroxide in basic solution is eliminated by boiling, which converts hydrogen peroxide to water and oxygen gas. e. Excess stannous chloride in hot HCl is eliminated by boiling, which converts the stannous cation to the stannic cation.

17. Which statement is FALSE for an oxidation involving potassium dichromate? a. The dichromate ion is orange in acidic solution. b. The dichromate ion is reduced to the chromous ion in acidic solution. c. Dichromate is a less powerful oxidizing agent than permanganate and cerium(IV). d. The dichromate ion is converted into the yellow chromate ion in basic solution. e. Potassium dichromate is a primary standard.

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Chap 16_10e 18. Cerium(IV) is a strong oxidizing agent commonly used in redox titrations. Which of the following is FALSE for cerium(IV)? a. Ce4+ is yellow and Ce3+ is colorless. The color change is abrupt enough that Ce4+ can serve as its own indicator. b. A Ce4+ solution prepared from ammonium hexanitratocerate(IV) can be used directly without standardization. c. Ce4+ does not form aqueous complexes when dissolved in water. d. Ce4+ is indefinitely stable in sulfuric acid solution. e. Ce4+ reacts rapidly with chloride in hot hydrochloric acid solutions to form chlorine gas.

19. Starch is an indicator for iodine. For iodimetry, the starch solution is added to the a. analyte solution at the beginning of the titration. b. analyte solution at the halfway point of the titration. c. analyte solution right before the equivalence point of the titration. d. analyte solution after the equivalence point of the titration has been reached. e. titrant solution before starting the titration.

20. Sometimes the oxidation state of the analyte needs to be adjusted before it is titrated. The analyte can be prereduced with a reducing agent to quantitatively reduce the analyte. The excess reducing agent is then eliminated to prevent interference with the titration. Which of the reducing agents is paired with an INCORRECT elimination chemistry? a. Stannous chloride in hot HCl is eliminated by the addition of excess HgCl2, which converts the stannous cation to the stannic cation. b. Chromous chloride is eliminated when excess chromium(II) reacts with atmospheric oxygen to form chromium(III). c. Sulfur dioxide is eliminated by boiling in acidic solution. d. Hydrogen sulfide is eliminated by boiling in basic solution. e. Excess hydrogen peroxide in acidic solution is eliminated by boiling, which converts hydrogen peroxide to water and oxygen gas.

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Chap 16_10e 21. Extra virgin olive oil is legally permitted to contain 20 mmol/kg of peroxides. Hydroperoxides in the presence of KI undergo the following redox reaction: The iodine release is titrated using sodium thiosulfate with a starch indicator until the blue color disappears and the solution turns colorless.

Calculate the amount of hydroperoxides (in mmol/kg) in an extra virgin olive oil sample if during the analysis, a chemist used 2.55 mL of 0.001 00 M sodium thiosulfate to titrate 20.00 mg of the oil sample. Does the olive oil contain less than the legally permitted amount of peroxides?

22. Consider the titration of 100.00 mL of 0.010 0 M Tl+ with 0.010 0 M in 1.00 M HCl, using a Pt indicator electrode and a saturated calomel reference electrode (E(saturated KCl) = 0.241 V).

Calculate the potential at the equivalence point.

E° = 1.24 V

Tl3+ + 2e−

Tl+

Eo = 0.77 V

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Chap 16_10e 23. The maximum level of ethanol (C2H5OH, FM 46.07)H5OH, FM that can be accurately detected by evidential breath testers is 220 μg/100 mL air. The redox reaction used by evidential breath testers is

Calculate the volume of 0.100 0 M K2Cr2O7 necessary to react with 220 μg of ethanol in a 100 mL sample of air.

24. The percent purity of a malonic acid reagent needs to be determined. A student weighs out 0.151 1 g of malonic acid (CH2(CO2H)2, FM 104.062) and quantitatively transfers the solid to a 250-mL Erlenmeyer flask. To the flask, 50.00 mL of distilled water is added to dissolve the malonic acid. After dissolution, the student uses a 50-mL volumetric pipet to add 50.00 mL of 0.200 M Ce4+, which is allowed to react with the malonic acid. After complete reaction, the excess Ce4+ is back titrated with 14.58 mL of 0.100 M Fe2+.

Calculate the percent purity of the malonic acid.

Ce4+ + Fe2+

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Ce3+ + Fe3+

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Chap 16_10e 25. Consider the titration of 50.00 mL of 0.020 0 M Cu+ by 0.025 0 M Fe3+ using a Pt indicator electrode and a saturated calomel reference electrode (E(saturated KCl) = 0.241 V). Calculate the cell potential after 20.00 mL, 40.00 mL, and 80.00 mL of Fe3+ has been added.

Fe3+ + e− Fe2+

Eo = 0.771 V

Cu2+ + e− Cu+

Eo = 0.161 V

26. A 10.0-g sample containing MnCl2 (FM 125.844) is taken for analysis. H2O2 is used to preoxidize Mn2+ to . The excess of H2O2 is eliminated by boiling, converting the hydrogen peroxide to water and oxygen gas. The is titrated with 15.75 mL of 0.012 5 M oxalic acid (H2C2O4) in 1 M H2SO4. Calculate the weight percent of MnCl2 in the sample.

27. A student is trying to determine the mass of iron in a women’s multivitamin. He takes four vitamin tablets and grinds them into a fine powder. He then reduces all iron in the sample to the +2 oxidation state. The resulting Fe2+ solution is transferred to a 250-mL volumetric flask and diluted to volume with distilled water. A 50.00mL aliquot of the sample requires 42.98 mL of 1.00 × 10−3 M

. Calculate the milligrams of Fe per

tablet.

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Chap 16_10e 28. A 50.00-mL aliquot of a water sample containing nitrite is reacted with excess potassium iodide in acidic solution to generate . Carbon dioxide is bubbled through the solution to remove any nitrogen monoxide that is generated. The water sample is transferred to a 500-mL volumetric flask and diluted to volume. A 50.00mL aliquot is then titrated against 1.092 × 10−4 M thiosulfate, requiring 15.48 mL to reach the starch end point. What is the

concentration in ppm? Assume a solution density of 1.000 g/mL.

29. A permanganate solution is prepared by dissolving 20.012 3 g KMnO4 (FM 158.034) in 500 mL of distilled water and boiling for 1 hour to remove any organic material. Following sintered-glass filtration, the solution is quantitatively transferred to a 1.00-L volumetric flask and diluted to volume with distilled water. The permanganate solution was titrated against 0.102 3 M oxalic acid prepared in sulfuric acid solution. A 50.00mL aliquot of oxalic acid solution required 16.68 mL of permanganate solution. A titration blank required 0.04 mL of permanganate. What is the permanganate molarity?

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Chap 16_10e 30. Consider the titration of 50.00 mL of 0.100 M Co3+ with 0.110 M . Calculate the potential after 50.00 mL of 0.110 M has been added. The potential is measured against the saturated Ag|AgCl reference electrode (E(saturated KCl) = 0.197 V).

Co3+ + e−

Co2+

+ e−

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Eo = 1.92 V

Eo = 0.56 V

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Chap 16_10e Answer Key 1. a 2. a 3. d 4. b 5. a 6. d 7. d 8. c 9. e 10. c 11. e 12. c 13. a 14. d 15. e 16. e 17. b 18. a 19. a 20. d 21. 63.75 mmol/kg; no; Use the concentration and volume of the thiosulfate solution to determine the number of millimoles added. Then use the reaction stoichiometries to determine the millimoles of hydroperoxides. Use the mass of the oil sample to calculate the amount in mmol/kg.

22. 0.842 V;

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Chap 16_10e

23. 31.84 μL; Use the mass of ethanol and its molar mass to determine the number of moles. Use the mole-to-mole ratio to calculate the moles of dichromate. Then use the concentration to determine the volume.

24. 98.0% pure malonic acid; Calculate the millimoles of cerium(IV) added and the millimoles of cerium(IV) in excess.

The difference is the millimoles of cerium(IV) that react with malonic acid. Using the mole-to-mole ratio, calculate the millimoles of malonic acid. Convert to moles and then to grams with the molar mass. Next, divide the mass of malonic acid by the mass of the sample and multiply by 100 to determine the percent purity of the malonic acid.

25. at 20.00 mL, E = −0.080 V; at 40.00 mL, E = 0.225 V; at 80.00 mL, E = 0.530 V; The equivalence point volume

is 40.00 mL. At 20.0 mL, [Cu2+] = [Cu+], so E = 0.161 V − 0.241 V = −0.080 V. At 40.00 mL, E = (0.771 V

+ 0.161 V)/2 – 0.241 V = 0.225 V. At 80.0 mL, [Fe2+] = [Fe3+], so E = 0.771 V − 0.241 V = 0.530 V.

26. 0.099 1%; Use the concentration and volume of the oxalic acid solution to calculate the moles of oxalic acid. Then use the 2:5 mole-to-mole ratio to determine the moles of permanganate, which are equal to the moles of MnCl2.

Use the molar mass of MnCl2 to determine the mass. Divide by the mass of the sample and multiply by 100 to determine the weight percent.

27. 18.0 mg Fe/tablet; Use the volume and concentration of the dichromate solution to calculate the millimoles of dichromate reacted. Use the mole-to-mole ratio to obtain the millimoles of iron reacted. Next, convert millimoles Fe to milligrams Fe using the molar mass of Fe. Divide the milligrams of Fe by 50.00 mL to get the mg Fe/mL for the original sample volume (250 mL). Multiply by 250 mL to find the total milligrams of Fe in the 250-mL solution.

Divide by four tablets to get 18.0 mg Fe/tablet.

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Chap 16_10e 28. 15.55 ppm ; The acidic solution protonates the nitrite, and the excess iodide reacts with nitrite to form

. The

percent yield is assumed to be 100% due to excess iodide. The sample is diluted to 500 mL and a 50.00-mL aliquot is titrated against thiosulfate to give the millimoles of in the 50.00-mL aliquot. The millimoles of nitrate in the aliquot can be determined from the millimoles of

. The concentration of nitrite in the 500-mL solution is

3.380 × 10−4 M. Assuming a 1.000 g/mL density for the water sample, the ppm nitrite is calculated as 15.55 ppm.

29. 0.123 0 M; Use the volume and concentration of the oxalic acid solution to get millimoles of oxalic acid. Then use the 2:5 mole-to-mole ratio to determine the millimoles of permanganate reacted. Subtract the titration blank volume from the permanganate volume to get the true volume of permanganate. Divide the millimoles by the volume in milliliters to determine the concentration. If you were to incorrectly use the mass of KMnO4—instead of the titration data—to determine the concentration, you would calculate a concentration of 0.1266 M. 30. 0.422 V; Use the volume and concentration of each species to determine that the titrant is in excess. Calculate the potential using the Nernst equation with 0.50 mmol excess MnO42− and 5.00 mmol MnO4−, E = 0.619 V. Subtract the reference electrode potential to determine the cell potential.

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Chap 17_10e Indicate the answer choice that best completes the statement or answers the question.

1. The _________ is the electrode at which the reaction of interest occurs. a. auxiliary electrode b. counter electrode c. anode d. cathode e. working electrode 2. Electrolysis of 200 mL of a solution containing an unknown metal ion increased the cathode mass by 0.043 4 g when a current of 0.500 A was passed through the cell for 5.00 min. The oxidation state of the metal ion is +2. What is the identity of the metal ion? a. Cr2+ b. Ni2+ c. Fe2+ d. Cd2+ e. Zn2+ 3. During electrolysis, electrons flow from the ____________ terminal of the power supply into the ___________ of the electrolysis cell. Electrons flow from the _____________ into the ________________ terminal of the power supply to complete the circuit. a. negative; anode; ground; positive b. positive; cathode; anode; negative c. negative; ground; cathode; positive d. negative; cathode; anode; positive e. positive; anode; cathode; negative 4. A student has five solutions containing Co2+ at various concentrations: (A) 0.10 M, (B) 0.25 M, (C) 0.40 M, (D) 0.50 M, and (E) 0.60 M. She performs electrolysis on 10.00 mL of one of the solutions. A constant current of 2.57 A is passed through the solution for 5.00 min. Which solution did the student perform the electrolysis on? a. solution (A) b. solution (B) c. solution (C) d. solution (D) e. solution (E)

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Chap 17_10e 5. Which statements regarding polarography are TRUE? I. Polarography is amperometry conducted with a dropping Hg electrode.

II. The diffusion current in the plateau region of the polarographic wave is proportional to the concentration of the analyte and is used for quantitative analysis.

III. The half-wave potential for maximum current is characteristic of a given analyte in a given medium and can be used for qualitative analysis for the analyte.

IV. Fresh Hg drops give reproducible results.

V. Polarography is used to primarily study oxidations. Hg may reduce when used to study reductions. a. I and V b. II, III, and V c. I, III, and IV d. III, IV, and V e. II, III, and IV 6. Which statement regarding electrogravimetric analysis is FALSE? a. Analyte is quantitatively deposited on an electrode via electrolysis. b. The increase in the mass of the electrode correlates to the amount of analyte in solution. c. If the analyte forms a colored solution, the disappearance of the color indicates completion of the deposition. d. If a fresh electrode surface is exposed to the analyte solution and further deposition does not occur, the deposition of the analyte is complete. e. If a qualitative test for a small sample of the analyte solution tests positive, the deposition of the analyte is complete.

7. A current of 0.500 A flows through a cell containing Fe2+ for 10.0 min. Calculate the maximum number of moles of Fe that can be removed from solution. Assume the current is constant over time. a. 1.04 mmol b. 51.8 μmol c. 3.11 mmol d. 1.55 mmol e. 25.9 μmol

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Chap 17_10e 8. A chemist wants to remove 99.99% of Cu2+ from a solution containing Cu2+, Sn2+, Fe2+, and Pd2+, each with a concentration of 0.10 M, without removing any of the other ions.

Cu2+ + 2e− ⇌ Cu(s) E° = 0.339 V

Sn2+ + 2e− ⇌ Sn(s) E° = −0.141 V

Fe2+ + 2e− ⇌ Fe(s)

E° = −0.440 V

Pd2+ + 2e− ⇌ Pd(s) E° = 0.915 V Which ion(s) will begin to reduce before 99.99% of the Cu2+ has been reduced to Cu(s)? a. Pd2+ b. Fe2+ c. None of the ions will be reduced before 99.99% of the Cu2+ is reduced. d. Fe2+, Pd2+ e. Sn2+

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Chap 17_10e 9. Which definition(s) is/are INCORRECT? I. Ohmic potential is the voltage to overcome the electrical resistance of the solution in an electrochemical cell when a current is flowing.

II. Concentration polarization occurs when the concentrations of the reactants or products at the surface of an electrode are greater than those in the bulk solution.

III. Overpotential is the voltage required to overcome the activation energy for a reaction at only the cathode.

IV. The voltage needed to drive a reaction is defined as E = E(cathode) – E(anode) – IR – overpotentials – concentration polarization. a. II, III, and IV b. I, II, and III c. III d. II and IV e. I and II 10. Which statement about square wave voltammetry, stripping analysis, and cyclic voltammetry is FALSE? a. Square wave voltammetry applies a square waveform superimposed on a staircase waveform of increasing potential to the working electrode. b. Stripping analysis is the most sensitive of the voltammetric techniques. c. Cyclic voltammetry applies a triangular-shaped waveform to the working electrode. d. For cathodic stripping analysis, the peak current during oxidation is proportional to analyte concentration. e. In cyclic voltammetry, for a reversible reaction, the peak current Ipc for the forward sweep of the first cycle is proportional to the concentration of analyte and the square root of the sweep rate.

11. The role of the supporting electrolyte in an unstirred analyte solution in cyclic voltammetry is to ______________ of the analyte to the electrode surface. a. increase diffusion b. reduce diffusion c. reduce migration d. increase migration e. reduce migration and increase diffusion

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Chap 17_10e 12. Consider the following electrolysis reactions. ⇌

Cathode: Anode:

⇌

The cell has a resistance of 65.0 Ω and a current of 4.75 mA when the cathode overpotential is 0.346 V and the anode overpotential is 0.062 V. Calculate the voltage needed to overcome these effects and drive the net reaction. The standard reduction potential for the half-reaction at the cathode is E° = 0.34 V and the standard reduction potential of the half-reaction at the anode is E° = 1.23 V. a. −1.578 V b. −1.607 V c. −0.860 V d. −1.169 V e. −1.268 V 13. Which statement about overpotential is FALSE? a. Overpotential is the voltage required to overcome the activation energy for a reaction at an electrode. b. The rate of the reaction will increase as the applied overpotential increases. c. Activation energy, and hence overpotential, is electrode-surface independent. d. Applying an increasing overpotential will sustain a higher current density. e. Current density is current per unit area of electrode surface, A/m2.

14. The half-wave potential in a voltammogram for a solution containing a redox couple that is being oxidized or reduced at a rotating disk electrode is equal to: a. the formal potential at which [red] = [ox]. b. the formal potential at which [red] = ½[ox]. c. the formal potential at which ½[red] = [ox]. d. half the formal potential at which [red] = ½[ox]. e. half the formal potential at which ½[red] = [ox].

15. For a three-electrode cell, the working electrode is the electrode at which the reaction of interest occurs.

The _______________ is used to measure the potential of the working electrode, and the _________________ is the current-supporting partner of the working electrode. a. auxiliary electrode; reference electrode b. reference electrode; auxiliary electrode c. reference electrode; polarizable electrode d. cathode; anode e. potential electrode; reference electrode

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Chap 17_10e 16. Which of the following statements are TRUE for coulometry?

I. Coulometric methods count electrons used in a chemical reaction to measure the amount of analyte.

II. Coulometry employs either constant-current or controlled-potential conditions.

III. An advantage of coulometric methods is high precision, and a disadvantage is low sensitivity.

IV. A coulometric titration is a controlled-potential method.

V. Controlled-potential coulometry utilizing a three-electrode cell is more selective than constantcurrent coulometry. a. I, II, and V b. I, II, IV, and V c. I, III, and IV d. II, IV, and V e. I and II 17. The Karl Fischer titration measures traces of water in samples, such as transformer oils, foods, and polymers. Which of the following is/are FALSE regarding the chemistry involved in a Karl Fischer titration?

I. The anode solution contains I−, an alcohol (ROH), a base (B), sulfur dioxide, and possibly other organic compounds.

II. III. I2 is generated from I− at the cathode.

IV. V. Two moles of electrons correspond to one mole of H2O. a. III b. I and III c. I, III, and V d. II and IV e. V 18. __________ measures the electric current flowing between a pair of electrodes driving an electrolysis reaction. The amount of current is proportional to the concentration of the analyte. a. Square wave voltammetry b. Polarography c. Cyclic voltammetry d. Amperometry e. Coulometry

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Chap 17_10e 19. Molecules have three ways to reach the surface of an electrode: diffusion, convection, and migration.

Which of these three methods control(s) the flux of analyte to the surface of a rotating disk electrode? a. convection b. diffusion and convection c. migration d. convection and migration e. diffusion and migration 20. _____________ is a set of techniques that observes the relationship between current and voltage during an electrochemical process. a. Amperometry b. Coulometry c. Electrophoresis d. Voltammetry e. Electrogravimetric analysis 21. A current of 0.250 A is passed for 30.0 min though a Cu2+ aqueous solution using a disk electrode of 3.00-mm diameter. The density of the deposited Cu(s) (FM 63.55) film is 8.96 g/cm3. Calculate the thickness of the deposited film.

22. Sulfide in a water sample is titrated with electrolytically generated Fe2+. If a 10.00-mL aliquot of the water sample requires a current of 0.750 A for 5.40 min to reach the end point, what is the molar concentration of sulfide in the water sample?

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Chap 17_10e 23. The Cl− concentration of a solution was determined by coulometric titration. Ag+ generated at the anode reacts with the Cl− in solution, forming the precipitate AgCl(s). The end point is determined when an increase in current is detected due to excess Ag+ in solution. A current of 0.500 A was applied for 17.55 min through a 25.00-mL solution containing an unknown Cl− concentration. Calculate the Cl− concentration.

24. An experiment measured the cathodic current for five solutions of known concentration using cyclic voltammetry. The collected data are plotted, and a best-fit line is fitted to the data. An unknown ferricyanide solution has a cathodic current of 17.38 mA. Calculate the concentration of ferricyanide in the unknown.

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Chap 17_10e 25. A technician is tasked with testing the accuracy of a newly purchased chloride autotitrator. He prepares a solution that is 3.423 8 mg NaCl/mL solution and takes a 25.00-mL aliquot for titration. To minimize any problems with the in situ generation of the Ag+ titrant, a fresh silver anode is installed, and the instrument is set to titrate with a constant current of 500.0 mA. The instrument took 286.3 s to reach the end point. Calculate the percent error for the instrument (NaCl FM 58.44).

26. When a potential is applied to a solution that is 0.10 M Ni2+ and 0.20 M Fe2+, which metal will be reduced first and at what potential? Is it possible to remove the metal from solution with 99.99% efficiency before the second metal begins to be reduced from solution?

Ni2+ + 2e− Ni(s) Eo = −0.236 V Fe2+ + 2e− Fe(s) Eo = −0.44 V

27. Based on a cyclic voltammetry experiment, it was determined that Epc = 0.329 V and Epa = 0.387 V versus S.H.E. for the reduction of ferricyanide to ferrocyanide at a scan rate of 10 mV/s. Determine if the reaction is reversible or quasi-reversible.

28. Hydrogen and oxygen gas are generated when current is passed through a solution of sodium sulfate.

Calculate the potential required for the cell to generate partial pressures of 0.25 bar for oxygen and 0.50 bar for hydrogen under the following conditions. At the cathode surface, the pH is 10.00 and at the anode surface, the pH is 4.00. A current of 0.32 A is passed through the solution, which has a resistance of 3.0 W. The overpotentials for the anode and cathode are 0.580 V and 0.584 V, respectively.

O2 + 4H+ + 4e− 2H2O Eo = 1.229 1 V 2H2O + 2e−

H2 + 2OH− Eo = −0.828 0 V

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Chap 17_10e 29. An experiment measured the cathodic current for five solutions of known concentration using cyclic voltammetry at a scan rate of 100 mV/s. The current was plotted versus the concentration, and the slope of the best-fit line was 7.45 × 104 A/(mol/L). The diffusion coefficient of is 3.09 × 10−6 cm2/s.

Calculate the geometric surface area of the electrode at which the cathodic process took place.

30. A current of 3.56 A is passed through an Fe(NO3)2 (FM 179.855) solution for 1.60 h. How many grams of Fe(s) (FM 55.845) will be deposited?

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Chap 17_10e Answer Key 1. e 2. c 3. d 4. c 5. e 6. e 7. d 8. a 9. a 10. d 11. c 12. a 13. c 14. a 15. b 16. a 17. a 18. d 19. b 20. d 21. 2.34 mm; Use the current, time, Faraday constant, stoichiometry of the reduction, and molar mass of iron to determine the mass of iron deposited. Use the density to determine the volume of the film. Assuming the film is cylindrical, calculate the height of the film.

22. 0.126 M S2−; Multiply the current by time expressed in seconds to get the total coulombs passing through the cell.

Next, use the Faraday constant to calculate the moles of electrons passing through the cell. The stoichiometry is 2 mol e− to 1 mol S2−. mol S2−/0.010 00 L = [S2−] 23. 0.218 3 M; Use the current, time, and Faraday constant to determine the number of moles of Ag+ generated, which is equivalent to the number of moles of Cl− in solution. Then use the solution volume to calculate the concentration.

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Chap 17_10e 24. 7.22 mM; Use the trendline data to calculate the concentration ferricyanide for 17.38 mA cathodic current.

25. 1.296%; There are two methods that can be used. The first method uses the titrator time, the amperage, and the volume of the sample to calculate the mg/mL concentration of NaCl and compares it to the known concentration.

The second method uses the known concentration, the amperage, and the volume of the sample to calculate the amount of time it should take to titrate the sample. This time is compared to the actual time.

26. A potential of −0.2656 V is required to start reducing Ni2+ from solution. A potential of −0.3839 V is required to remove 99.99% of Ni2+ from solution. A potential of −0.4607 V is required for reduction of Fe2+ from solution.

Therefore, it is possible to remove Ni2+ from the solution before Fe2+ is reduced. Use the Nernst equation to calculate the potential required when the concentration is 0.10 M Ni2+, −0.2656 V. Next, calculate the potential required to start reducing Fe2+ from solution, −0.4607 V. Comparison of the reduction potentials indicates that Ni2+ will reduce first. Calculate the potential when 99.99% of Ni was removed from solution (0.01% still in solution), −0.3839 V. The Fe2+ potential is too negative to start reduction at −0.3839 V. 27. Reversible;

E° = 0.358 V S.H.E. 0.387 V – 0.329 V = 0.058 V/1 = 58 mV

28. −3.690 1 V; First, determine which half-reaction occurs at the cathode and which occurs at the anode. Since this is an electrolytic cell, the more negative potential remains the reduction and is the cathode. The more positive electrode is the anode. Insert the desired partial pressure of hydrogen gas and the hydroxide concentration into the Nernst equation for the cathode half-reaction. Insert the desired partial pressure of oxygen and the hydronium concentration into the Nernst equation for the anode half-reaction. Solve for the potential of the cell (E = Ecathode – Eanode), which includes the concentration polarization at the electrode surfaces. Then subtract the ohmic resistance (IR) and the overpotentials for the anode and cathode.

29. 0.501 cm2; Use the equation

, where the slope of the line is equal to Ipc/C, to

calculate A. 30. 5.93 g; Use the time, current, and Faraday constant to determine the number of moles of electron passed through the solution. Then use the stoichiometry of the reduction to determine the number of moles of Fe2+ reduced.

Finally, use the molar mass of Fe to determine the mass of Fe(s) deposited.

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Chap 18_10e Indicate the answer choice that best completes the statement or answers the question.

1. Which statements related to the Beer-Lambert law are correct?

I. Absorbance increases as concentration increases.

II. Absorbance decreases as pathlength increases.

III. The molar absorptivity value is chromophore and wavelength specific.

IV. The absorption spectrum is a graph showing how absorbance changes with wavelength. a. I, II, and III b. I, III, and IV c. II, III, and IV d. I and III e. I, II, III, and IV 2. Which statement is NOT true for light modeled as a wave? a. Frequency is the number of oscillations that a wave makes each second. b. Wavelength is the distance between the crest of one wave and the trough of the next wave. c. Light travels slower through matter than through a vacuum. d. The speed of light in a vacuum is 2.998 × 108 m/s. e. The product of the frequency and wavelength of a wave of light is the speed of light.

3. The excited state in which the electron spins are opposed is called a. the triplet state. b. the singlet state. c. the ground state. d. phosphorescence. e. fluorescence.

4. Beer’s law states that absorbance is proportional to the concentration of the absorbing species. Which of these does NOT result in a deviation from this ideal relationship? a. stray light b. high analyte concentrations c. instrumental electrical noise d. polychromatic radiation e. high ionic strength of analyte solution

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Chap 18_10e 5. The part of a spectrophotometer that selects the wavelength of light used to irradiate the sample is called the a. polychromator. b. monochromator. c. beam splitter. d. light source. e. optode.

6. Reagent blanks are primarily used to a. align the optics to achieve maximum irradiance. b. compensate for any absorbance due to reagents or contamination. c. verify reagent purity. d. calibrate the spectrophotometer. e. test detector sensitivity.

7. Phosphorescence is a. a nonradiative transition between states with the same spin, S1→ S0. b. a nonradiative transition between states with different spin, T1→ S0. c. the emission of a photon during a transition between states with the same spin, S1→ S0. d. the emission of a photon during a transition between states with different spin, T1→ S0. e. None of these statements describe phosphorescence.

8. Which statement about luminescence is FALSE? a. Luminescence is emission of light from an excited state of a molecule. b. Luminescence is inherently more sensitive than absorption. c. At low concentration, emission intensity is proportional to analyte concentration. d. Emission comes at higher energy (shorter wavelengths) than absorption. e. Luminescence is quenched by self-absorption at high analyte concentrations.

9. Calculate the energy of a photon with a frequency of 573.8 nm. a. 3.46 × 10−37 J b. 3.80 × 10−40 J c. 3.46 × 10−28 J d. 3.80 × 10−31 J e. 3.46 × 10−19 J

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Chap 18_10e 10. Which statement is TRUE for an emission spectrum? a. The excitation wavelength is fixed, and the emission wavelengths are scanned. b. The excitation wavelength is scanned, and the emission wavelength is fixed. c. The excitation wavelength is fixed, and the emission wavelength is fixed. d. The excitation wavelength is scanned, and the emission wavelengths are scanned. e. None of these statements describe how an emission spectrum is collected.

11. Which statement is TRUE for an excitation spectrum? a. The excitation wavelength is fixed, and the emission wavelengths are scanned. b. The excitation wavelength is scanned, and the emission wavelength is fixed. c. The excitation wavelength is fixed, and the emission wavelength is fixed. d. The excitation wavelength is scanned, and the emission wavelength is scanned. e. None of these statements describe how an excitation spectrum is collected.

12. Absorbance measured during a spectrophotometric titration must be corrected due to a. changes in the absorbing wavelength over time. b. dilution of the analyte solution by the addition of the titrant. c. contaminants in the titrant. d. instrumental drift. e. decomposition of analyte.

13. The basic spectrophotometer is composed of four components. Which component is NOT one of the four components in a basic spectrophotometer? a. sample cell b. monochromator c. beam splitter d. light source e. light detector 14. A 10.00-mL sample is titrated with 0.50 mL of titrant. The observed absorbance is 0.321 9. Calculate the corrected absorbance. a. 0.306 6 b. 0.305 8 c. 0.291 2 d. 0.338 0 e. 0.355 8

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Chap 18_10e 15. Which statement is NOT true for light modeled as a particle? a. A particle of light is the photon. b. The energy of a photon is proportional to its frequency. c. The wavenumber of a photon is inversely proportional to its energy. d. The energy of a photon decreases as its wavelength increases. e. The product of frequency and Planck’s constant equals photon energy.

16. Which of the following statements is NOT true about the absorption of light? a. Transmittance is the fraction of incident light that is not absorbed. b. When a molecule absorbs light, it is promoted to an excited state. c. A = −log T d. When a molecule emits light, it returns to a lower energy state. e. Irradiance is the energy per second of the light beam.

17. What material should be used for a cuvette if a technician needs to measure a sample that absorbs at 255 nm in the ultraviolet region of the electromagnetic spectrum? a. KBr b. quartz c. glass d. fused silica e. polystyrene 18. If 99% of the light is absorbed by a solution, then according to the equation

a. A = 1. b. A = 2. c. A = 0.1. d. A = 0.01. e. A = 0.099.

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Chap 18_10e 19. A 3.70 × 10−4 M solution of compound X has an absorbance of 0.534 2 when measured at 740 nm.

Assuming a pathlength of 1.00 cm, calculate the molar absorptivity of compound X at 740 nm. a. 1.44 × 103 M−1 cm−1 b. 9.93 × 105 M−1 cm−1 c. 5.06 × 103 M−1 cm−1 d. 736 M−1 cm−1 e. 1.44 × 105 M−1 cm−1 20. ________ is the energy per unit time per unit area in the light beam. a. Irradiance b. Absorbance c. Transmittance d. Molar absorptivity e. Frequency 21. Which statement about the excited and ground states of molecules is FALSE? a. When a molecule absorbs light of sufficient energy, it is much more likely to undergo the transition S0 → T1 than the transition S0 → S1. b. Electron spins are parallel in the triplet, T1, excited state and opposed in the singlet, S1, excited state. c. Internal conversion is a nonradiative transition from the S1 state to the S0 state. d. When a molecule absorbs light of sufficient energy, an electron transitions from the S0 to the S1 state. e. Intersystem crossing is a nonradiative transition from the T1 state to the S0 state.

22. Beer’s law works for _________ radiation passing through a _________ solution in which the absorbing species is _________ in a concentration-dependent equilibrium. a. polychromatic; dilute; not participating b. monochromatic; concentrated; not participating c. monochromatic; concentrated; participating d. polychromatic; dilute; participating e. monochromatic; dilute; not participating 23. Fluorescence is a. a nonradiative transition between states with the same spin, S1 → S0. b. a nonradiative transition between states with different spin, T1→ S0. c. the emission of a photon during a transition between states with the same spin, S1 → S0. d. the emission of a photon during a transition between states with different spin, T1→ S0. e. None of these statements describe fluorescence.

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Chap 18_10e 24. A sample with an analyte concentration of 1 × 10-6 M produces an emission intensity (I) of 0.40 and has an absorbance (A) of 0.090. What will be the emission intensity and absorbance of the sample if the incident irradiance, P0, is doubled?

25. A transmittance of 0.384 7 is measured for a 2.300 × 10−4 M solution. Assuming a pathlength of 1.000 cm, calculate the molar absorptivity of the solute.

26. The iron content of a well water sample must be quantified. A 10.00-mL aliquot of well water is transferred to a 100.0-mL volumetric flask, and 20 mL of 0.10 M nitric acid is added to oxidize iron(II) to iron(III).

Next, 30 mL of a 0.1 M sodium thiocyanate solution is added to form the deep red complex iron(III) thiocyanate. The solution is diluted to volume with distilled deionized water. A reagent blank is similarly prepared. If the absorbance of the iron thiocyanate solution is 0.548 3 and the reagent blank absorbance is 0.028 2, calculate the iron concentration for the water sample. The molar absorptivity for iron(III) thiocyanate is 3 200 M−1 cm−1. Assume a pathlength of 1.000 cm and a 100% yield.

Fe3+(aq) + SCN-(aq) ↔ FeSCN2+(aq)

27. Quinine sulfate in 0.5 M sulfuric acid is excited at 310 nm and fluoresces at 450 nm. How many joules are dissipated through nonradiative transitions on a per molecule basis?

28. After 10.00 mL of titrant has been added to 50.00 mL of the analyte solution, the absorbance was measured to be 0.463 9. Calculate the corrected absorbance.

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Chap 18_10e 29. A solution of the iron-transport protein transferrin was titrated with a standard solution of Fe3+. The blank absorbance was 0.04, and the absorbance of 10.00 mL of the transferrin solution measured after adding 6.00 mL of titrant was 0.65. Calculate the corrected absorbance that should be plotted.

30. A red laser emits a pulse of light every 10.0 ms. If the laser operates at a wavelength of 695 nm and 50.0 W, calculate the number of photons emitted during a pulse. (1 W = 1 J/s)

31. Calculate the wavelength of light necessary to promote an electron from the valence band to the conduction band within a semiconductor nanoparticle with a bandgap of 3.02 eV.

32. The absorbance of a sample and blank measured in a 1.00-cm cuvette was 0.057 and 0.003, respectively.

Calculate the absorbance of the sample if it was measured in a 10.00-cm cuvette.

33. A dye solution of unknown concentration has an emission intensity of 0.346 in fluorescence analysis. Then 0.20 mL of a 1.56 µg/mL solution of the dye was mixed with 9.00 mL of the unknown solution and diluted to 10.00 mL in a volumetric flask. The emission intensity of the solution is 0.584. Find the concentration of the dye in the unknown sample.

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Chap 18_10e Answer Key 1. b 2. b 3. b 4. e 5. b 6. b 7. d 8. d 9. e 10. a 11. b 12. b 13. c 14. d 15. c 16. e 17. b 18. b 19. a 20. a 21. a 22. e 23. c 24. I = 0.80; A = 0.090; When the irradiance is doubled, the emission intensity will double because I = kP0c. In contrast, doubling the irradiance has no effect on the absorbance because it is equal to the ratio of the irradiance exiting and entering the sample.

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Chap 18_10e 25. 1 804 M−1 cm−1; Calculate the absorbance from the transmittance (0.414 9). Then use Beer’s law to calculate the molar absorptivity.

26. 1.625 × 10-3 M; Calculate the absorbance of the sample corrected for the blank (0.520 1). Then use Beer’s law to calculate the iron concentration (0.000 1625 M) of the diluted sample. Finally, calculate the concentration of iron in the water sample, accounting for the dilution factor (1:10).

27. 1.99 × 10−19 J dissipated; Calculate E310 and E450, and subtract E450 from E310 to determine the energy dissipated, Edisp. Or solve for x in the following equation, which is the wavelength of the energy difference. x = 996.4 nm. Then calculate the energy difference from the wavelength, 1.99 × 10−19 J. 28. 0.556 7; Divide the total volume of the solution by the initial volume of the analyte solution and multiply by the observed absorbance.

29. Corrected A = (16.00 mL/10.00 mL)(0.65 – 0.04) = 0.98 30. 1.75 × 1015 photons; Calculate the energy of one photon (2.86 × 10–19 J/photon). Then, calculate the total energy of one pulse (0.000 500 J). Finally, divide the total energy by the energy of a photon (1.75 × 1015 photons).

31. 411 nm; 1 eV = 1.602 18 × 10-19 J; 3.02 eV × 1.602 18 × 10-19 J/1 eV = 4.838 58 × 10-19 J; ΔE = hn = hc/l; 4.838 58 × 10-19 J = (6.626 × 10-34 J∙s)(2.998 × 108 m/s)/l; l = 4.11 × 10-9 m = 411 nm 32. A = 0.057(10.00 cm) – 0.003(10.00 cm) = 0.54 33. 0.039 6 µg/mL; Standard addition problem; [dye]f = [dye]i (9.00 mL/10.00 mL); [S]f = (1.56 µg/mL)(0.20 ml/10.00 mL) = 0.031 2 µg/mL; [dye]i/(0.031 2 µg/mL + 0.900[dye]i) = 0.346/0.584; [dye]i = 0.039 6 µg/mL

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Chap 19_10e Indicate the answer choice that best completes the statement or answers the question.

1. The degree of photobleaching is dependent on all of the following EXCEPT a. the propensity of the fluorophore to populate the triplet state. b. the number of times the fluorophore is excited. c. the lifetime of the excited state. d. the intensity of the excitation light and duration of exposure. e. the reversibility of the photobleaching process.

2. ____________ is the emission of light from a living system, and _____________ is the emission of light from a chemical reaction. a. Chemiluminescence; bioluminescence b. Luminescence; fluorescence c. Bioluminescence; chemiluminescence d. Bioluminescence; luminescence e. Fluorescence; chemiluminescence 3. Which of the following is TRUE for the isosbestic point for a reaction in which absorbing species X is converted to absorbing species Y? a. [X] = [Y]

b. εX + εY = constant and [X] + [Y] = constant

c. εX = εY and [X] = [Y] d. εX + εY = constant and [X] = [Y] e. εX = εY and [X] + [Y] = constant 4. Which of the following are advantages of sequential injection over flow injection analysis?

I. Less waste is generated.

II. Smaller volumes of expensive reagents are consumed.

III. It provides the ability to handle a larger number of samples in the same period of time.

IV. It provides the ability to be miniaturized and used in remote locations. a. I, II, and IV b. III and IV c. II and IV d. I, II, III, and IV e. I and II

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Chap 19_10e 5. A chemist uses gel filtration chromatography to purify a protein. She collects 40 fractions from the column and needs to determine which fractions contain the protein. She only has 500 μL of a 0.250-mg/mL standard solution of the protein to prepare a calibration curve. Which technique is best suited for this application? a. 96-well microplate b. colorimeter c. discrete analyzer d. flow injection analysis e. sequential injection 6. The Stern-Volmer equation relates quencher concentration [Q] to the ratio Φ0/ΦQ. Which statement regarding the Stern-Volmer equation is INCORRECT? a. The y-intercept for the Stern-Volmer equation is always 1. b. The x-intercept for the Stern-Volmer equation is always 1. c. The slope for the graph of Φ0/ΦQ versus [Q] is

.

d. I0/IQ is proportional to Φ0/ΦQ. e. Φ0/ΦQ is greater than 1 for any concentration of Q. 7. Sequential injection differs from flow injection analysis by a. flow programming. b. flow reversal. c. reagent mixing. d. flow reversal and reagent mixing. e. flow programming and flow reversal.

8. _________ enhances the fluorescence of immunoassays by a factor of 100 by measuring the fluorescence of Eu3+ 200 μs after excitation with a laser pulse. a. Time-resolved fluorescence immunoassays b. Time-dependent fluorescence immunoassays c. Immunoassay fluorescence over time immunoassays d. Weakly bound time fluorescence immunoassays e. Europium-bound fluorescence immunoassays 9. Which statement regarding spectroscopy measurements based on scattering by particles is INCORRECT? a. Turbidimetry measures the amount of light that passes through the sample in the forward direction. b. Nephelometry measures the amount of light that is scattered at an angle by the sample. c. Turbidimetry is well suited for samples with a high concentration of suspended particles. d. Nephelometry is well suited for samples with a low concentration of suspended particles. e. Sample solutions can be colored or colorless.

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Chap 19_10e 10. Which statements about ELISA are TRUE? I. An antibody, antibody 1, specific to the analyte of interest is dissolved in aqueous solution.

II. Analyte is incubated with antibody 1, binding to form a complex. The fraction of antibody 1 sites bound to analyte is proportional to the concentration.

III. Unbound substances are flushed from the ELISA surface and a second antibody, antibody 2, which binds to a different region of the analyte, is applied to the antibody 1-analyte complexes.

IV. Antibody 2 is chemically modified to include an enzyme that converts a colorless reagent to a colored reagent. The more bound analyte, the more intense the color generated. a. I, II, III, and IV b. II and IV c. I, III, and IV d. II, III, and IV e. I, II, and IV 11. Which of the following are advantages of flow injection analysis over discrete analyzers?

I. the speed of sample analysis II. automation of solution handling III. reproducibility IV. low cost of analysis a. II and III b. I, II, III, and IV c. I, III, and IV d. II, III, and IV e. I, II, and III 12. When a molecule absorbs a photon, the molecule is promoted to an excited state, M*. The rate at which M* is created is proportional to the concentration of M. There are three pathways that M* may take to return to the ground state. Which statement regarding these three pathways is INCORRECT? a. The rate at which the excited state returns to the ground state is independent of the pathway to the ground state and the concentration of excited state molecules. b. The three possible pathways to return to the ground state are emission, deactivation, and quenching. c. Quenching occurs when a second molecule, the quencher Q, collides with M*, which returns M* to the ground state. M* + Q → M + Q* d. Deactivation returns M* to the ground state by colliding with other molecules and releasing energy in the form of heat. M* → M + heat e. Emission returns M* to the ground state by emitting a photon. M* → M + hν

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Chap 19_10e 13. The technique in which a liquid sample is injected into a continuously flowing liquid carrier containing a reagent that reacts with the sample is called a. flow analysis. b. sequential flow analysis. c. sequential analysis. d. sequential injection analysis. e. flow injection analysis.

14. Which statement regarding static fluorescence quenching is FALSE? a. Quencher and fluorophore form a stable nonfluorescent ground-state complex. b. Fluorescence intensity depends on the concentration of the quencher. c. Fluorescence lifetime is independent of the quencher concentration. d. Fluorescence lifetime depends on the quencher concentration as described by the Stern-Volmer equation. e. The lifetime of the fluorescing species is the time needed for the luminescence intensity to fall to 37% of its initial value.

15. Most compounds are not luminescent. To analyze such compounds using fluorescence detection, a fluorescent moiety is coupled to the compound through a process called a. modification. b. derivatization. c. quenching. d. photobleaching. e. fluoroactivation.

16. In fluorescence spectroscopy, the quantum yield is defined as a. the rate of fluorescence emission. b. the number of photons emitted. c. the number of photons emitted divided by the number of photons absorbed. d. the number of excitation photons impinging on a sample divided by the number of photons absorbed. e. the fraction of excited molecules produced by direct excitation.

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Chap 19_10e 17. A plot of absorbance versus wavelength is shown for two different compounds. Which two wavelengths are most suited for the analysis of a mixture of the two compounds?

a. 515 nm, 515 nm b. 445 nm, 635 nm c. 380 nm, 750 nm d. 515 nm, 635 nm e. 445 nm, 515 nm 18. Which statement is TRUE for the relationship between quencher concentration and fluorescence intensity? a. As quencher concentration increases, fluorescent intensity increases. b. As quencher concentration decreases, fluorescent intensity decreases. c. As quencher concentration increases, fluorescent intensity decreases. d. As quencher concentration increases, fluorescent intensity remains constant. e. As quencher concentration decreases, fluorescent concentration remains constant.

19. Which statement is TRUE for the spectroscopic measurement of a mixture? a. The absorbance of a mixture at a particular wavelength is the sum of the absorbances of the components in the mixture that absorb at the particular wavelength. b. Spectrophotometers can differentiate between components of a mixture that absorb at the same wavelength. c. Each component in a mixture has the same molar absorptivity at the same wavelength. d. The concentration for each component in the mixture is easily calculated with a set of simultaneous equations when there is significant overlap of the individual spectra. e. The concentration for each component in a mixture is easily calculated by least-squares analysis when the individual spectra are well resolved.

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Chap 19_10e 20. Which statements about Fӧrster resonance energy transfer are FALSE? I. It is a nonradiative energy transfer due to dipole-dipole interactions between adjacent molecules that are touching.

II. The emission spectrum of the donor must overlap with the emission spectrum of the acceptor for energy transfer to occur.

III. Energy goes from a donor in the ground state to an acceptor in the excited state.

IV. The likelihood of energy transfer decreases with the sixth power of distance between the donor and acceptor. a. I and III b. II and III c. I, II, and III d. I, III, and IV e. III and IV 21. An acid-base indicator has the following equilibrium: HIn ⇌ In− + H+. The absorbance data in the table were collected at pH 1 and pH 13 using 5.00 × 10−5 M solutions of HIn and In−. The absorbance of a solution containing a mixture of the two forms of the indicator was collected at pH 7. Calculate the equilibrium constant (K) for the equilibrium reaction at pH 7. A 1.00-cm cell was used to measure the absorbance.

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Chap 19_10e 22. Consider the following reaction: Fe3+ + SCN− FeSCN2+. A series of FeSCN2+ standards was prepared and the absorbance was measured at 450 nm. The data in the table was used to prepare a calibration curve.

Absorbance [FeSCN2+] (M) 0 0 0.190 4.00 × 10−5 0.370 8.00 × 10−5 0.560 1.20 × 10−4 0.770 1.60 × 10−4

A 5.00-mL aliquot of a 2.00 × 10−3 M Fe(NO3)3 solution was mixed with a 4.00-mL aliquot of a 1.00 × 10−3 M KSCN solution and 1.00 mL of HNO3. The absorbance of this mixture at equilibrium was 0.71.

Calculate the equilibrium constant for the reaction.

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Chap 19_10e 23. A microplate is used to determine the DNA concentration of several hundred samples by measuring the absorbance at 260 nm. Each microwell has a volume of 250 μL. The absorbance of a DNA sample in one microwell, the absorbance of a blank in a separate microwell, and the absorbance of water in a 1.000-cm quartz cuvette are shown in the table.

Calculate the pathlength of the DNA sample and use it to determine the concentration of the DNA sample in μg/mL. The absorptivity of DNA at 260 nm is 0.0284 (μg/mL)−1cm−1.

24. For the acid-base indicator bromothymol blue, the protonated form (HB) is yellow and the deprotonated form (B−) is blue. When both forms are present the indicator has a green color. The shade of green depends on the ratio of HB:B−. Use the information in the table to calculate the concentrations of HB and B − for a green solution of bromothymol blue. A 1.00-cm cell was used to measure the absorbance.

Wavelength Absorbance of Absorbance of Absorbance of −5 −5 − Unknown Mixture 1.00 × 10 M HB 1.00 × 10 M B 453 nm 0.191 0 0.023 4 0.468 0 616 nm 0.006 25 0.321 1 0.326 9

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Chap 19_10e 25. An oxygen sensor is constructed from a fluorescent dye that is quenched by oxygen. The sensor is calibrated with a deoxygenated water sample and two water samples with known concentrations of dissolved oxygen. The measured fluorescence intensity is provided in the table along with the Stern-Volmer plot. The sensor is deployed to a waste treatment plant to monitor dissolved oxygen in effluent. If the sensor records a fluorescence intensity of 322.928, what is the oxygen concentration?

O2 (M) Fluorescence Intensity 0.00 537.283 372.992 1.30 × 10−4 294.482 2.60 × 10−4

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Chap 19_10e 26. In microbiology, microbial growth is monitored using turbidity at 600 nm. Calculate the optical density of a solution in which the absorbance due to scattering is 0.544 in a 0.500-cm cell.

27. A chemical reaction converts species X to species Y. The absorbance spectra of pure X and pure Y shows an isosbestic point at 650 nm. The absorbance measured at 650 nm for a mixture of the two species is 0.650 in a 1.000-cm cell. The concentrations of the species in the mixture are [X] = 2.73 × 10−4 M and [Y] = 4.72 × 10−4 M. Calculate the molar absorptivity of both species εX and εY at 650 nm.

28. The fluorescence of quinine is quenched by the chloride ion. The fluorescence intensity of quinine was measured in the presence of six different standard chloride solutions and the results were used to prepare a Stern-Volmer plot. The chloride standards were prepared using a 0.010 0 M standard chloride solution, adding an increasing volume of chloride to a series of 100-mL volumetric flasks. To each flask, 10.00 mL of 1.000 × 10−5 M quinine was added. Each solution was diluted to volume with 1 M H2SO4. For the unknown, 10.00 mL of the sample and 10.00 mL of 1.000 × 10−5 M quinine solution were added to a 100-mL volumetric flask and diluted to volume with 1 M H2SO4. Calculate the chloride concentration for the unknown.

Fluorescence Intensity [Cl−] (M) 0.000

140.00 137.70

1.000 × 10−3 134.70 2.000 × 10−3

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Chap 19_10e

[Cl−] (M) 0.000

Fluorescence Intensity 140.00 131.30

3.000 × 10−3 127.61 4.000 × 10−3 123.70 5.000 × 10−3 Unknown

133.50

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Chap 19_10e 29. From the UV-Vis absorption data in the table, calculate the concentrations of the metal ions M2+ and N2+ in a river sample. The unknown solution was prepared by mixing 10.00 mL of river water with 10.00 mL of ligand solution and diluting to a final volume of 100.00 mL. Assume the absorbance was measured in a 1.00-cm cell.

Solution Wavelength Concentration Absorbance 650 nm 0.723 M2+ standard 1.82 × 10−5 650 nm 0.312 N2+ standard 1.82 × 10−5 425 nm 0.199 M2+ standard 1.82 × 10−5 425 nm 0.811 N2+ standard 1.82 × 10−5 Unknown 650 nm ?

0.753 Unknown 425 nm ?

0.536

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Chap 19_10e 30. The fluorescence of quinine is quenched by the bromide ion. The lifetime of the quinine fluorescence was measured in the presence of several different concentrations of the bromide ion. The results were used to prepare a Stern-Volmer plot. The lifetime was also measured for an unknown bromide concentration.

Calculate the concentration of the unknown bromide solution. Is this an example of static or dynamic quenching?

Lifetime (ns) [Br−] (M) 18.1 8.9 5.7 3.6 2.5 1.9 2.8

0.000 0.005 0.010 0.020 0.030 0.040 unknown

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Chap 19_10e Answer Key 1. e 2. c 3. e 4. a 5. a 6. b 7. e 8. a 9. e 10. d 11. b 12. a 13. e 14. d 15. b 16. c 17. b 18. c 19. a 20. a

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Chap 19_10e 21. K = 1.95 × 10−8; At pH = 1, the indicator exists almost exclusively as HIn. At pH = 13, the indicator exists almost exclusively as In−. Use the absorbances of the 5.00 × 10−5 M solutions of HIn and In− to determine the

molar absorptivities at each wavelength. Set up two equations relating the absorbance for the unknown mixture at each wavelength to the molar absorptivities and unknown concentrations of the two species. Solve the equations simultaneously. Use the concentrations and the H+ concentration to calculate K. 22. 271; Use the calibration curve to calculate the equilibrium concentration of FeSCN2+. Set up an equilibrium table to determine the equilibrium concentrations of Fe3+ and SCN− from their initial concentrations and the equilibrium concentration of FeSCN2+. Insert these concentrations into the K expression and solve for K.

23. b = 0.600 cm; c = 30.5 μg/mL; Determine the pathlength of the solution in the well using the equation

. Then calculate the concentration from the corrected absorbance, A

= εbc.

24. 2.33 × 10−5 M HB and 9.73 × 10−6 M B−; Calculate the molar absorptivity for each species at each wavelength from the known concentrations for HB and B− and the absorbances. Set up two equations relating the absorbance for the unknown mixture at each wavelength to the molar absorptivities and unknown concentrations of the two species. Solve the equations simultaneously.

25. 2.06 × 10−4 M; Divide the fluorescence intensity of the deoxygenated sample by the fluorescence intensity of the unknown, insert it into the best-fit line equation for y, and solve for x, which is equal to the oxygen concentration.

26. 1.088; OD = A(due to scatter)/(b/1.000 cm) = 0.544/(0.500 cm/1.000 cm) = 1.088

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Chap 19_10e

27. 872 M−1 cm−1; At the isosbestic point, e X = e Y, so the absorbance of the solution can be expressed as A =

εb([X] + [Y]).

28. 1.886 × 10−3 M Cl−; Divide the fluorescence intensity of the 0 M chloride standard by the fluorescence intensity of the unknown to obtain I0/IQ for the sample. Insert this value in for y in the best-fit line equation and solve for x, which is the chloride concentration in the dilute unknown. Multiply by 10 to determine the chloride concentration of the unknown.

29. [M2+] = 1.54 × 10−5 M and [N2+] = 8.25 × 10−6 M; Calculate the molar absorptivity for each ion at both wavelengths from the standard solutions. Set up two equations for the sample mixture, one each for the absorbance at 650 and 425 nm, and solve them simultaneously for [M2+] and [N2+].

30. 0.026 M; dynamic quenching; Divide the lifetime of the blank by the lifetime of the unknown, insert it into the best-fit line equation for y, and solve for x, which is equal to the bromide concentration of the unknown. This is an example of dynamic quenching because the lifetime of the fluorescence of the analyte is dependent on the concentration of the quencher.

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Chap 20_10e Indicate the answer choice that best completes the statement or answers the question.

1. Which statement about lasers is FALSE? a. Stimulated emission describes the phenomenon where a photon of a specific energy stimulates an excited molecule to emit a photon and return to the ground state. b. A laser consists of a lasing medium between two mirrors: one which reflects all light (0% transmittance) and one that reflects most light (1% transmittance). c. A necessary condition for lasing is population inversion, in which a higher energy state has a smaller population than a lower energy state in the lasing medium. d. Lasers are line sources that provide radiation of a single wavelength. Lasers are available that provide radiation from ultraviolet to infrared wavelengths. e. The round-trip distance of the optical cavity must be an exact multiple of the wavelength of the emitted radiation.

2. Spectrophotometers employ various sources of radiation. Which statement regarding spectrophotometer light sources is FALSE? a. A quartz-tungsten halogen lamp emits radiation over the visible spectrum and parts of the ultraviolet and infrared spectrums. b. A deuterium arc lamp uses controlled electric discharges to cause D2 to dissociate and emit visible radiation. c. Light-emitting diodes emit narrow bands of visible and near-infrared radiation. d. Lasers emit radiation of a single wavelength. e. A silicon carbide globar emits infrared radiation when heated to near 1 500 K. 3. Which source CANNOT be used as a source of ultraviolet or visible radiation? a. deuterium arc lamp b. xenon arc lamp c. mercury arc lamp d. silicon carbide globar e. quartz-tungsten halogen lamp 4. A(n) ____________ response is a function of incident light. An electrical current is produced when struck by ____________. For a given radiant power, the current produced is ______________ analyte concentration. a. optical; photons; constant for varying b. detector; photons; proportional to c. analyte; analyte molecules; constant for varying d. detector; analyte molecules; inversely proportional to e. optical; analyte molecules; inversely proportional to

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Chap 20_10e 5. Which statement regarding photodiode arrays is FALSE? a. Photodiode arrays respond to more than one wavelength. b. Photodiode arrays require a polychromator to detect more than one wavelength. c. Photodiode arrays allow for rapid acquisition of an entire light spectrum. d. Photodiode arrays have higher signal-to-noise ratios. e. Spectrophotometers using photodiode arrays are dispersive spectrometers.

6. Assuming a constant waveguide length and angle of incidence, why does the sensitivity of the attenuated total reflectance sensor increase with decreasing waveguide thickness? a. The sensitivity increases with decreasing waveguide thickness due to an increase in the reflection angle. An increased reflection angle causes the evanescent wave to penetrate farther into the sample, thereby increasing the attenuation of the laser beam and greater sensitivity. b. A thinner waveguide allows more light to pass through, resulting in a greater signal-to-noise ratio and increased sensitivity. c. The sensitivity of an attenuated total reflectance sensor increases with decreasing waveguide thickness due to a reduction in the angle required for total internal reflection. With a reduced reflection angle, the evanescent wave will penetrate to a greater depth for a greater period of time. This results in greater attenuation and greater sensitivity. d. The number of internal reflections increases as the waveguide thickness is decreased, resulting in a greater amount of attenuation and an increase in sensitivity. e. The sensitivity of the attenuated total reflectance sensor is independent of the waveguide thickness.

The sensitivity can only be changed by changing the waveguide length.

7. The signal-to-noise ratio of a spectrum is 4 after 100 spectra have been averaged. How many spectra would need to be averaged to produce a signal-to-noise ratio of 8? a. 200 b. 300 c. 400 d. 600 e. 800 8. Which monochromator-related term is NOT correctly defined? a. grating – optical element with closely spaced lines b. diffraction – bending of light by a grating c. refraction – scattering of light by a lens or prism d. polychromatic – multiple wavelengths e. monochromatic – one wavelength

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Chap 20_10e 9. Calculate the minimum grating length to successfully resolve 616.50 nm from 616.59 nm for a first-order diffraction. The available diffraction grating has 103 grooves per cm. a. 13.7 cm b. 1.00 cm c. 111.1 cm d. 0.15 cm e. 6.85 cm 10. Which property of laser light is NOT correctly defined? a. monochromatic – emits one wavelength of light b. extremely bright – high power at one wavelength c. collimated – parallel rays of light d. polarized – electric field oscillates between two perpendicular planes e. coherent – all waves in phase 11. Which statements are TRUE for optical fibers?

I. Optical fibers carry light by total internal refraction.

II. Optical fibers are constructed of a high-refractive-index transparent core enclosed in a lower refractive-index transparent cladding.

III. Optical fibers exploit Snell’s law, n1 sin θ1 = n2 sin θ2, to transmit light along the length of the optical fiber.

IV. If n1/n2 < 1, there is an allowed range of angles θi in which essentially all light is reflected at the walls of the core, with negligible amounts of light penetrating the cladding. a. I and IV b. I, II, and IV c. II, III, and IV d. II and III e. II and IV

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Chap 20_10e 12. Two components of a Michelson interferometer are a moveable mirror and a beam splitter. Which statements regarding these components are FALSE? I. When the mirror travels a distance of λ/4, deconstructive interference occurs.

II. The beam splitter alternates the incident light between a reference and the sample.

III. When the mirror travels a distance of λ/2, deconstructive interference occurs.

IV. The beam splitter splits the incident light between the stationary mirror and the moveable mirror. a. I and II b. I and IV c. II and III d. II and IV e. III and IV 13. Stray light reaching the detector is a problem for spectrophotometers that are not properly sealed. What is the impact of stray light on the recorded absorbance? a. The apparent absorbance is greater than the true absorbance and the difference between the two increases as the amount of stray light increases. b. The apparent absorbance is greater than the true absorbance and the difference between the two decreases as the amount of stray light increases. c. The apparent absorbance is less than the true absorbance and the difference between the two increases as the amount of stray light decreases. d. The apparent absorbance is less than the true absorbance and the difference between the two decreases as the amount of stray light decreases. e. It is not possible to determine the effect of stray light on the measured absorbance.

14. Which property of noise does signal averaging exploit to improve the signal-to-noise ratio? a. Noise is loud. b. Noise is random. c. Noise is predictable. d. Noise drifts. e. The amplitude of noise is a constant.

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Chap 20_10e 15. Which statements are TRUE for a photomultiplier tube?

I. A photomultiplier tube is a very sensitive detector, amplifying the photoelectron current by approximately 106.

II. Emitted photoelectrons are accelerated toward positively charged electrodes called dynodes.

III. Electrons strike the dynode with a kinetic energy less than the original emitted kinetic energy.

IV. For each electron striking a dynode, more than one electron is knocked free. The additional electrons are accelerated toward the next dynode in the detector, where amplification is repeated. a. I and II b. I and IV c. II, III, and IV d. I, II, and IV e. II and III 16. Which detector is NOT an appropriate detector for a spectrophotometer? a. photomultiplier tube b. thermal conductivity c. photodiode array d. charge coupled device e. phototube 17. A ____________ disperses light into its component wavelengths and selects a narrow band of wavelengths to pass on to the sample or detector. a. polychromator b. photodiode array c. monochromator d. charge coupled device e. photomultiplier tube 18. Which type of noise is NOT correctly defined? a. line (interference) noise – noise that occurs at discrete frequencies b. 1/ƒ (drift) noise – low-frequency noise caused by flickering or drifting of light intensity c. Gaussian noise – noise whose amplitude is dependent on frequency d. Johnson noise – random fluctuations of electrons in electronic devices e. shot noise – noise due to the quantized nature of charge carriers and photons

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Chap 20_10e 19. Which methods can be used to increase the resolution of a grating?

I Increase the total number of grooves by increasing the grating length II Optimize the diffraction order by choosing the correct blaze angle III Decrease the groove density while keeping the grating length constant IV Increase the size of the individual grooves while keeping the grating length constant V Match the blaze angle to the entrance slit width for maximum throughput a. II and III b. III and IV c. I and II d. I and V e. IV and V 20. Which statement regarding double-beam scanning spectrophotometers is FALSE? a. The reference beam allows for correction of source intensity fluctuations. b. Comparison of the irradiance emerging from the sample and reference cuvettes is used to determine absorbance and transmittance. c. The reference beam allows for correction of detector drift. d. The light beam is alternatively passed through the sample and reference cuvettes using a stationary mirror. e. A rotating mirror is used to direct the light beam that passes through the sample and reference cuvettes to the detector.

21. The signal-to-noise ratio for a single spectrum is 2.18. How many additional scans are required to increase the signal-to-noise ratio by a factor of 50?

22. Find the minimum angle of incidence (θi) that will result in total internal reflection for an optical fiber with a core that has a refractive index of 1.70 and a cladding that has a refractive index of 1.47.

23. By what factor will the exitance from a blackbody increase if the temperature of the blackbody object is tripled?

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Chap 20_10e 24. The dispersion of a diffraction grating (Δϕ/Δλ) is a measure of the angular separation of adjacent wavelengths. Calculate the dispersion of a first-order diffraction for light emerging at ϕ = 23° when a grating with 1 800 grooves per centimeter is used.

25. Ice has a refractive index of 1.309. What is the speed of light in ice?

26. Light travels from a medium with a refractive index of 1.567 into a second medium with a refractive index of 1.121 at an angle of 43°. At what angle does the light pass through the second medium?

27. The monochromator for a new spectrophotometer must be able to resolve lines that are 0.005 nm apart at 631 nm. Calculate the minimum number of grooves per centimeter required for the diffraction grating if the grating has a length of 5.00 cm. Assume first-order diffraction.

28. A spectroscopist recorded a very weak XPS signal coming from a sample containing a few atoms of Au deposited on a Ti surface. The signal-to-noise ratio was 1.6. How many spectra should the spectroscopist collect and average to improve the signal-to-noise ratio to 32?

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Chap 20_10e 29. The true absorbance for a 1.0 × 10−5 M solution is 0.752 6. If the percentage stray light for a spectrophotometer is 0.56%, calculate the percentage by which the apparent concentration deviates from the known concentration.

30. Calculate the first-, second-, and third-order wavelengths of light diffracted from a grating with 2 000 lines/mm (blazes), if the incident and reflection angles are 30° and 45°, respectively.

31. A Fourier transform infrared spectrophotometer contains a mirror that travels ±0.9 cm. Calculate the maximum retardation (Δ) and resolution of the instrument. For proper coverage of a spectral range (Δν) of 211 to 3200 cm−1, at what retardation interval (δ), in millimeters, does the interferogram need to be sampled?

32. An atomic emission spectrophotometer is used to monitor the emission from Ni at 345.85 nm and 346.17 nm. Can the monochromator in the spectrophotometer successfully resolve 345.85 nm from 346.17 nm?

The grating is 5.00 cm in length and there are 5 500 grooves per cm. Assume first-order diffraction.

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Chap 20_10e 33. A portion of the experimental data from a noisy X-ray fluorescence (XRF) spectrum, where the signal in counts per second is plotted on the y-axis as a function of the X-ray energy in kiloelectron volts on the xaxis, is shown in the table.

X-ray energy Signal X-ray energy Signal X-ray energy Signal (keV) (counts/s) (keV) (counts/s) (keV) (counts/s) 10.925 1.430 11.125 99.454 11.325 124.319 10.950 2.860 11.150 71.488 11.350 102.042 10.975 12.425 11.175 87.015 11.375 24.864 11.000 15.527 11.200 205.128 11.400 22.133 11.025 1.430 11.225 202.025 11.425 34.186 11.050 24.906 11.250 198.923 11.450 23.720 11.075 12.425 11.275 167.840 11.475 9.322 11.100 27.966 11.300 132.568 Use the seven-point centered moving average smoothing algorithm and the seven-point Savitzky-Golay polynomial smoothing procedure to calculate the smoothed value of the ordinate, y, for an abscissa of x = 11.375 keV.

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Chap 20_10e Answer Key 1. c 2. b 3. d 4. b 5. e 6. d 7. c 8. c 9. e 10. d 11. d 12. c 13. c 14. b 15. d 16. b 17. c 18. c 19. c 20. d 21. A total of 2 500 total scans are required, so 2 499 additional scans. The signal-to-noise ratio increases in proportion to , where n is the number of scans.

22. 59.8°; Use Snell’s law to determine the angle (θi) at which

= 1, which is when no light is passed into the

cladding and all of the light is reflected internally.

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Chap 20_10e 24. 11.2°/µm; Determine the distance between the grooves (d) in micrometers per groove from the number of grooves per centimeter. Then calculate Δϕ/Δλ in radians per micrometer using the equation Δϕ/Δλ = 1/(d cos ϕ). Convert the answer to degrees per micrometer.

25. 2.29 × 108 m/s; The speed of light in a medium of refractive index n is c/n, where c is the speed of light in a vacuum.

26. 72°; Use Snell’s law to calculate θ2.

27. 25 240 grooves/cm; Insert λ, Δλ, and n into the resolving power of a grating equation and solve for N, the number of required illuminated grooves. Divide N by the length of the grating to determine the number of grooves per centimeter required.

28. 400; The signal-to-noise ratio increases in proportion to

, where n is the number of scans.

29. 1.48%; Calculate the molar absorptivity from the known absorbance and concentration. Next, calculate the apparent absorbance from the true absorbance and the percent stray light. Use the apparent absorbance and the molar absorptivity to calculate the apparent concentration. Then calculate the percent deviation using the equation ((true concentration – apparent concentration)/true concentration) ∙100.

30. first-order: 603 nm; second-order: 301.5 nm, third-order: 201 nm; Insert the values for the order (n); distance between grooves (d), which is the inverse of the blazes; incident angle (θ); and reflection angle (ϕ) into the grating equation and solve for the wavelength (λ).

31. Δ = 1.8 cm; resolution = 0.56 cm−1; δ = 0.001 7 mm; The maximum retardation is twice the distance of the mirror movement. The resolution is equal to 1/Δ. The retardation interval is given by the equation δ = 1/2Δν.

32. The resolving power of the monochromator is 25.40 times greater than needed. The required resolution to differentiate between the Ni emission lines is 345.85 nm/0.32 nm = 1 080.8. The resolving power of the monochromator is calculated by first determining Dl, 345.85 nm/Dl = 1(27 500 grooves); Dl = 0.012 6 nm. The resolving power is 345.85 nm/0.012 6 nm = 27 448.4.

33. Using the seven-point centered moving average method y = 66.262 counts/s. Using the seven-point Savitzky-Golay method y = 51.526 counts/s.

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Chap 21_10e Indicate the answer choice that best completes the statement or answers the question.

1. What percentage of Na atoms are in the lowest excited state in an acetylene–air flame at 2 500 K? The lowest excited state of sodium lies 3.371 × 10−19 J/atom above the ground state. The degeneracy of the excited state is 2, whereas the degeneracy of the ground state is 1. a. 0.011 4% b. 0.002 86% c. 99.989% d. 100.0% e. 0.00% 2. Which mechanisms are responsible for linewidth broadening in atomic spectroscopy?

I. pressure broadening II. ionization effect III. matrix effect IV. Doppler effect a. I and IV b. II, III, and IV c. I, III, and IV d. II and IV e. III and IV 3. Which technique is NOT used for background correction in atomic spectroscopy? a. beam chopping b. deuterium lamp background correction c. Zeeman background correction d. subtracting adjacent pixels of a CID display e. Fourier transformation 4. Which statement about ionization interference is FALSE? a. Ionization interference can occur with alkali metals at relatively low temperature and for analyses of other elements at higher temperature. b. Ionization suppressors decrease the extent of ionization of analyte. c. Ionization suppressors are more easily ionized than is the analyte. The large number of electrons produced when the suppressor is ionized prevent ionization of the analyte. d. Ionization interference occurs when a component of the sample prevents ionization of the analyte. e. All of the statements are true.

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Chap 21_10e 5. Which statements are TRUE for chemical interference?

I. Releasing agents are added to the sample to reduce chemical interference.

II. Chemical interference occurs when a component of the sample reduces the extent of atomization of the analyte.

III. Higher temperature flames will reduce chemical interference.

IV. Chemical interference occurs when a component of the sample emits light at or close to the same frequency as the analyte.

V. Releasing agents will preferentially react with the analyte, protecting the analyte from components of the sample that will ionize the analyte. a. III, IV, and V b. I, III, and IV c. I, II, and IV d. I, II, and III e. II, III, and V 6. The dynamic reaction cell reduces isobaric interference with thermodynamically favorable reactions that remove interfering species with the same mass-to-charge ratio. Which reaction does not remove interfering species for 52Cr+ or 56Fe2+? a. 40Ar12C+ + NH3 → + Ar + C b. 40Ar16O+ + NH3 →

+ Ar + O

c. 35Cl16OH+ + NH3 →

+ ClO

d. 40Ca16O+ + CO → Ca+ + CO2 e. 35Cl17OH+ + NH3 →

+ ClO

7. A(n) ___________ is a highly ionized, electrically neutral gaseous mixture of cations and electrons that approaches temperatures of 10 000 K, and is commonly used as an excitation source in atomic emission spectroscopy. a. inductively coupled plasma b. flame in a laminar flow burner c. electrothermal atomization d. spark e. arc

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Chap 21_10e 8. Laser ablation is a sampling technique used with inductively coupled plasma–mass spectrometry. Which of the statements regarding laser ablation are TRUE? I. Laser ablation destroys the entire sample.

II. A laser pulse is focused onto a small area, producing an explosion of atoms, electrons, and ions in the gas phase.

III. Depth profiling of the sample is possible with repeated laser pulses probing deeper and deeper into the sample.

IV. Nanogram amounts of material are ablated per pulse.

V. Ablation products are swept to the plasma with H2 gas. a. I and II b. II, III, and IV c. I, II, III, and IV d. II, IV, and V e. III and IV 9. Which statement is NOT true for X-ray fluorescence? a. The fluorescent X-ray photon has a lower energy than the original X-ray photon. b. The incoming X-ray knocks free an electron from the K, L, or M shells. c. The fluorescent X-ray photon is emitted when an outer shell electron falls into a shell closer to the nucleus. d. K X-rays originate when L or M shell electrons fall into the K shell. e. Kα and Kβ X-rays originate when L shell electrons transition to the K shell.

10. Which statements are true regarding the effect of flame temperature on atomic absorbance and atomic emission spectroscopy?

I. Varying the flame temperature by 10 K hardly affects the ground-state population of the analyte and would not noticeably affect the analyte signal in atomic absorption spectroscopy.

II. Varying the flame temperature by 10 K hardly affects the excited-state population of the analyte and would not noticeably affect the analyte signal in atomic emission spectroscopy.

III. Varying the flame temperature by 10 K affects the ground-state population of the analyte and will noticeably affect the analyte signal in atomic absorption spectroscopy.

IV. Varying the flame temperature by 10 K affects the excited-state population of the analyte and will noticeably affect the analyte signal in atomic emission spectroscopy. a. I and II b. I and IV c. II and III d. III and IV e. II and IV

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Chap 21_10e 11. What are the advantages to using a furnace instead of a flame in atomic absorption?

I. Sample atomization occurs in one heating step.

II. A smaller amount of sample is required.

III. Less operator skill is required to determine the proper experimental conditions.

IV. There is higher sensitivity because the atomized sample is in the optical path longer.

V. Memory effects from the previous run are not present. a. I, III, and V b. II, III, and IV c. II and IV d. I, II, and III e. II and V 12. __________ is any effect that changes the signal while the analyte concentration remains unchanged. a. Scattering b. Interference c. Self-absorption d. Diffraction e. Refraction 13. Inductively coupled argon plasma does not suffer from many of the interferences encountered with flame atomic spectroscopy. Which statement does NOT present an advantage of using plasma over a flame? a. Plasma is twice as hot as a conventional flame. b. Formation of analyte oxides and hydroxides is negligible. c. Atomization is more complete and the signal is enhanced in plasma. d. Analyte in plasma self-absorbs at a higher rate than in a flame. e. The analyte residence time in plasma is about twice as long as in a flame.

14. Atomic line broadening resulting from frequency fluctuations due to the motion of atoms in the flame or furnace is known as a. pressure broadening. b. uncertainty broadening. c. Zeeman broadening. d. Doppler broadening. e. Heisenberg broadening.

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Chap 21_10e 15. Which method is NOT used as a background correction technique in atomic spectroscopy? a. Adjacent pixel height is subtracted from signal pixel height for a charge injection device (CID) detector. b. A rotating beam chopper periodically blocks the analyte signal, allowing for a background measurement. c. The absorbance and scattering of the light from a deuterium lamp is measured and subtracted from the absorbance and scattering of light from a hollow-cathode tube. d. The analyte sample is spiked with a compound to absorb all background signal. e. A strong magnetic field is pulsed on and off. The sample and background are observed when the field is off. Only the background is observed when the field is on. The absorbance difference is the analyte signal.

16. Complete the statements regarding the effect the energy difference between the ground and excited states has on the excited-state population, and the effect the flame temperature has on the excited-state population.

At a given temperature, as the energy difference between the ground state and excited state decreases, the excited-state population ______________. As the temperature of the flame increases, the excited-state population _______________. a. remains the same; increases b. decreases; decreases c. increases; remains the same d. increases; increases e. decreases; remains the same 17. The natural linewidth of an atomic absorption line is governed by a. Doppler broadening. b. the Heisenberg uncertainty principle. c. pressure broadening. d. the Boltzmann distribution. e. nebulization.

18. A graphite furnace employs temperature programming to produce elemental vapor. The program begins by raising the furnace temperature to initiate ___________, where water is removed from the sample. Next, the temperature of the furnace is raised to begin the ___________ process, where organic material is removed. The temperature of the furnace is raised again to __________ the sample, creating the elemental vapor required for spectroscopy. a. baking; desiccation; atomize b. drying; desiccation; fragment c. drying; charring; atomize d. drying; atomize; charring e. baking; atomize; desiccation

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Chap 21_10e 19. There is no universal flame temperature for flame-based atomic spectroscopy that gives optimal results for all elements. _____________ gives a ____________. a. Excess oxidant; cooler flame b. Excess oxidant; hotter flame c. Shorter aspiration time; cooler flame d. Longer aspiration time; hotter flame e. Excess analyte; hotter flame 20. Which statements are TRUE of atomic emission spectroscopy (AES)?

I. Atoms are promoted to an excited state with a laser.

II. Atoms are promoted to an excited state by gaining energy from collisions with other atoms or from the high thermal energy of the flame.

III. The intensity of emitted light is proportional to analyte concentration.

IV. A plasma often replaces the flame in AES. a. III and IV b. I, III, and IV c. II, III, and IV d. II and III e. I and III 21. In X-ray fluorescence, X-rays are generated when: a. metals such as W, Mo, Ag, or Rh are heated to high temperatures. b. high-energy photons strike a metal such as W, Mo, Ag, or Rh. c. high-energy electrons strike an anode made of metals such as W, Mo, Ag, or Rh. d. high-energy protons strike an anode made of metals such as W, Mo, Ag, or Rh. e. metals such as W, Mo, Ag, or Rh are electrically excited.

22. A flame can be used in both atomic absorption and atomic emission spectroscopy to atomize the sample. In which technique is a stable flame temperature more critical? Why? a. A stable flame temperature is more critical in atomic absorption spectroscopy because a small fluctuation in the temperature can greatly affect the ground-state population of atoms in the sample. b. A stable flame temperature is more critical in atomic emission spectroscopy because a small fluctuation in the temperature can greatly affect the excited-state population of atoms in the sample. c. A stable flame temperature is more critical in atomic emission spectroscopy because a small fluctuation in the temperature can greatly affect the amount of sample that is atomized. d. A stable flame temperature is more critical in atomic absorption spectroscopy because a small fluctuation in the temperature can greatly affect the excited-state population of atoms in the sample. e. A stable flame temperature is equally important in atomic absorption and atomic emission spectroscopy because a small fluctuation in the temperature can greatly affect the amount of sample that is atomized.

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Chap 21_10e 23. _________________ occurs in inductively coupled plasma–mass spectroscopy when the interfering species has the same mass-to-charge ratio as the analyte ion. a. Ionization interference b. Analyte interference c. Isobaric interference d. Allotrope interference e. Isothermal interference 24. Which statement regarding hollow-cathode lamps is FALSE? a. The cathode of a hollow-cathode lamp is made of the same element whose emission lines are desired. b. The low-pressure Ne or Ar gas is ionized when voltage is applied, which ionizes the gas and accelerates the ions toward the cathode. c. The kinetic energy of the ionized gas is sufficiently high to sputter metal atoms from the cathode. d. The gaseous metal atoms are excited by collisions with high-energy electrons and emit photons with characteristic wavelengths that are absorbed by the analyte. e. The emission from the lamp is wider than the absorption line of the atoms in the flame.

25. Which statement is NOT true of the flame atomic absorbance spectroscopy of iron? a. The iron sample is aspirated into the flame, where the solvent evaporates and the iron analyte is atomized. b. The flame replaces the cuvette of conventional spectrophotometry, and the flame pathlength is typically 10 cm. c. The hollow-cathode lamp for iron emits wavelengths of light unique for iron. d. The iron atoms in the flame absorb some of the light emitted from the hollow-cathode lamp. e. A detector measures the amount of light generated by the flame and not absorbed by iron.

26. A sample analyzed by laser ablation–inductively coupled plasma–mass spectrometry has a density of 3.57 g/mL. After a 19.0-ns laser pulse, a crater with a diameter of 30.8 µm and a depth of 0.583 µm was created in the sample. If the energy of the laser pulse is 2.52 mJ, determine the laser power density in watts per square centimeter (W/cm2). How much sample mass was removed by the laser pulse? Assume the crater has a cylindrical shape.

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Chap 21_10e 27. A peak was observed in an X-ray fluorescence (XRF) spectrum at 13.115 keV. Calculate the wavelength of this X-ray radiation in angstroms.

28. Determine the Doppler linewidth of the 361.939-nm emission line of Ni at 2 940 K.

29. An element reaches its first excited state by absorption of 375.3 nm light. If the degeneracies of the two states for the element are g*/g0 = 3, determine N*/N0 at 2 010 K. By what percentage does N*/N0 change if the temperature is raised by 20 K?

30. A hair sample was analyzed for Pb using atomic absorption spectroscopy. The sample was treated with 5.00 mL of concentrated nitric acid and then diluted to 100.0 mL with deionized water. The sample solution was aspirated directly into the air–acetylene flame, yielding an absorbance of 0.485 at 283.3 nm. Pb standard solutions yielded the following absorbance values.

Concentration (ppm) Absorbance 0.250 0.386 0.500 0.589 Calculate the concentration of Pb in ppm in the hair sample.

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Chap 21_10e 31. The method of standard addition was used to measure the calcium content in a breakfast cereal. A 0.659 0g sample of crushed cereal was reacted for 1 h in warm 1 M HNO3. The reaction mixture was filtered and the filtrate was transferred to a 100-mL volumetric flask. A series of eight solutions were prepared by transferring 5.00-mL aliquots of the 100-mL solution to 50-mL volumetric flasks to which increasing volumes of standard Ca2+ (containing 20.0 mg/mL) was added, diluting to the final volume with deionized H2O. The samples were then analyzed by flame atomic absorption. Use the data in the table and the graph to find the percent by mass calcium in the cereal.

Ca2+ Std

Abs

(mL) 0

0.151

1.00

0.185

3.00

0.247

5.00

0.300

8.00

0.388

10.00

0.445

15.00

0.572

20.00

0.723

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Chap 21_10e 32. Calculate the energy released in kilojoules per mole, when a sample of gallium emits Kα radiation at 9.252 keV.

33. An element with energy levels E0 and E* has degeneracies for those energy levels of g0 and g*, respectively. Determine the fraction of atoms in the excited state at 3 321 K if the wavelength difference of the two states is 541.0 nm, g0 = 2, and g* = 3.

34. Calculate the emission wavelength for excited atoms that lie 3.055 × 10−19 J above the ground state.

35. Ba was used as an internal standard to analyze a sample containing Cr using atomic absorption spectroscopy (AAS). A standard mixture containing 2.04 µg Ba/mL and 1.79 µg Cr/mL measured by AAS produced a Ba signal/Cr signal = 2.58/1.00. A mixture was prepared by combining 4.50 mL of a Cr solution of unknown concentration with 4.00 mL of a 13.0-µg/mL solution of Ba. The absorbance of the mixture at the Cr wavelength was 0.104 and the absorbance at the Ba wavelength was 0.210. Determine the concentration of Cr in the unknown solution.

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Chap 21_10e Answer Key 1. a 2. a 3. e 4. d 5. d 6. e 7. a 8. b 9. e 10. b 11. c 12. b 13. d 14. d 15. d 16. d 17. b 18. c 19. b 20. c 21. c 22. b 23. c 24. e 25. e

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Chap 21_10e 26. 1.78 × 1010 W/cm2; 1.55 ng; Calculate the power of the laser pulse by dividing the energy by the time. Determine the power density by dividing the power by the area of the crater. Then determine the volume of the crater, assuming it is a cylinder. Use the volume and density to calculate the mass removed.

27. 0.945 36 Å; Convert the energy from kiloelectronvolts to joules. Then calculate the wavelength and convert it to angstroms.

28. 1.79 × 10−3 nm; Use the following equation to calculate the Doppler linewidth (δλ), where T is the temperature, λ is the emission wavelength, and M is the atomic mass.

29. 1.570 × 10−8; 20.7%; Calculate ΔE for the excited state and use it in the equation

to calculate

N*/N0, where k is the Boltzmann constant and T is the temperature. Repeat the calculation at the second temperature. Divide the change in N*/N0 by N*/N0 at the initial temperature and multiply by 100 to determine the percentage change.

30. 0.372 ppm; Use the absorbance values and concentrations of the standard solutions to determine the slope of the calibration line. Then use the slope, concentration, and absorbance of one of the standards to determine the yintercept. Use the equation of the line to calculate the concentration of the sample from the absorbance 31. 0.340 wt% Ca; The absolute value of the x-intercept is equal to the concentration of the sample in the 50-mL volumetric flask. Multiply by 50 mL to determine the mass of calcium in the flask. Then, multiply by the dilution factor (100 mL/5.00 mL) to determine the mass of calcium in the sample. Convert the mass of calcium to grams and solve for the percentage.

32. 8.927 × 105 kJ/mol; Convert the energy from kiloelectron volts per atom to kilojoules per atom. Use Avogadro’s number to convert the energy to kilojoules per mole.

33. 0.000 500; Calculate ΔE from the wavelength difference of the two states. Then use the Boltzmann distribution equation to calculate the fraction of atoms in the excited state.

34. 650.2 nm; Use the equation ΔE = hc/λ to solve for λ, where ΔE is the energy difference, h is Planck’s constant, and c is the speed of light in a vacuum.

35. 3.75 × 10−5 M; Use the standard solution mixture to determine the response factor F. Then use the response factor, the absorbance values, and the Ba concentration in the unknown mixture to calculate the concentration of the diluted sample. Use the dilution factor to calculate the concentration of the unknown solution.

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Chap 22_10e Indicate the answer choice that best completes the statement or answers the question.

1. Which structure best fits the mass spectrometry data?

a.

b.

c.

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Chap 22_10e

d. e. None of the structures fits with the mass spectrometry data.

2. Analytes on the surface of an object can be sampled using an open-air sampling technique by vaporizing and ionizing the analyte directly from the surface under atmospheric conditions. Which statements regarding these techniques are TRUE? I. In direct analysis in real time, excited He is generated, which produces protonated water clusters from distilled water. The water clusters protonate the analyte.

II. In low-temperature plasma, excited-state species are generated in a plasma, which ionize the analyte.

III. In desorption electrospray ionization, micron-sized droplets of analyte-free solvent created with electrospray are directed toward the surface of the object, dissolving the analyte in the droplets.

IV. All open-air sampling techniques dislodge analyte from the surface, and collect the ions for later mass spectrometer analysis. a. I and IV b. II and III c. I and III d. II and IV e. I and II

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Chap 22_10e 3. Which statements regarding the mass spectrometry of proteins are TRUE? I. MALDI and electrospray are major methods for introducing proteins into the gas phase.

II. The equation

is used to determine the number of protons bound to a protein from

the mass spectrum.

III. Once n is known and the molecular mass of the protein is calculated, the molecular mass (M) of the neutral protein is M = nmn .

IV. Electron-transfer dissociation is a mass spectrometry technique used to sequence proteins. a. I, II, and IV b. II and IV c. III and IV d. I, III, and IV e. II and III 4. Which statements are TRUE regarding the use of mass spectrometry as a chromatography detector?

I. Selected ion monitoring observes only a few m/z values to create a selected ion chromatogram.

II. An extracted ion chromatogram is created by recording the entire mass spectrum for the compounds coming off the column and picking one m/z value from the collected mass spectra.

III. A reconstructed total ion chromatogram is created by recording the total ion current below a certain m/z value.

IV. Selected reaction monitoring creates a chromatogram using a triple quadrupole mass spectrometer.

Quadrupole 1 selects the precursor ion. In quadrupole 2, the precursor ion is fragmented through collisions with N2 or Ar gas. All fragment ions pass to quadrupole 3, which selects one ion fragment to monitor. a. I, II, III, and IV b. II, III, and IV c. I and IV d. III and IV e. I, II, and IV

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Chap 22_10e 5. Which statement regarding mass spectrometers is FALSE? a. Quadrupole mass spectrometers use a conversion dynode followed by a continuous-dynode electron multiplier, in which ions striking the dynode eject electrons from the dynode toward the electron multiplier. b. Transmission quadrupole mass spectrometers select ions by applying a constant voltage and a radiofrequency oscillating voltage to the four parallel metal rods. c. Orbitrap mass spectrometers do not require magnetic or radio-frequency fields to measure m/z. d. Quadrupole mass spectrometers operate at constant resolving power. e. The time-of-flight mass spectrometer employs a reflectron to group ions with the same m/z, regardless of initial kinetic energy.

6. The resolving power for a mass spectrometer can be calculated using the following equations. or Which statements are NOT true regarding resolving power?

I. (m/z) is the larger value of m/z.

II. Δ(m/z) is the difference between the two m/z peaks.

III. (m/z)1/2 is the width of the peak at half maximum height.

IV. (m/z) is the smaller value of m/z.

V. (m/z)1/2 is the average mass for the two peaks. a. I and V b. II, III, and IV c. I, II, and III d. I, II, and V e. II and IV 7. How many rings and double bonds are there for the compound C11H7NO2Cl2? a. c = 10; h = 7; n = 1; R + DB = 8 b. c = 11; h = 9; n = 1; R + DB = 8 c. c = 10; h = 7; n = 3; R + DB = 9 d. c = 11; h = 7; n = 1; R + DB = 9 e. c = 11; h = 9; n = 3; R + DB = 9

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Chap 22_10e 8. Which statement regarding ion mobility spectrometry is FALSE? a. Ion mobility spectrometry is not a form of mass spectrometry. b. Ion mobility spectrometry is gas-phase electrophoresis. c. Ions are separated according to their size-to-charge ratio. d. Ions are generated when sample is desorbed and passes through a 10-mCi 63Ni source. e. All ions travel with the same kinetic energy.

9. The three essential components of any mass spectrometer are I. an ion source.

II. a decelerator.

III. a detector.

IV. a mass separator.

V. an accelerator. a. I, III, and IV b. II, III, and IV c. I, III, and III d. II, IV, and V e. III, IV, and IV 10. Analytes must be ionized prior to entering the mass filter in a mass spectrometer. Electron ionization and chemical ionization are two ionization techniques. Which statements regarding these ionization techniques are INCORRECT? I. Electron ionization uses an electron beam to create molecular ions, M+.

II. Chemical ionization uses an electron beam to ionize reagent gas.

III. Electron ionization gives very little fragmentation of the molecular ion.

IV. The ionized reagent gas in chemical ionization undergoes a complex set of chemical reactions before protonating the analyte to create MH+.

V. Both ionization techniques give identical mass spectra. a. II, IV, and VI b. I, II, and III c. I, III, and V d. I and III e. II, IV, and V

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Chap 22_10e 11. Which of the molecule(s) will have an odd-numbered molecular ion m/z value?

I.

II.

III.

IV.

V. a. II, III, and IV b. I and V c. I d. II and IV e. III and V

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Chap 22_10e 12. What is the purpose of the matrix in matrix-assisted laser desorption/ionization (MALDI)?

I. Upon irradiation, the matrix desorbs, carrying the analyte with it into the gas phase.

II. The matrix forms a complex with the analyte, making it easier to ionize.

III. The matrix is ionized by the laser and some of the charge is transferred to the analyte.

IV. It increases the fragmentation of the analyte.

V. A mixture of the analyte interspersed between fine crystals of the matrix is created once the solvent evaporates.

VI. It enhances the signal of low abundance or small molecule analytes. a. II, IV, and VI b. I, II, and III c. I, III, and V d. I and III e. II, IV, and V 13. Chemical ionization is one method that can be used to create ions from volatile analytes for analysis via mass spectrometry. Which statement describes how ions are created via chemical ionization? a. Analyte molecules are bombarded with high-energy electrons (usually 70 eV), allowing the molecules to absorb enough energy to ionize. b. A reagent gas is converted to reactive species by energetic electrons. These new reactive species then donate a proton to the analyte molecules to produce the protonated molecule, MH+. c. Analyte molecules are ejected from the solution phase to the gas phase. During this process the analyte molecules lose an electron, causing ionization. d. Analyte molecules are bombarded with protons that pull electrons away from the analyte molecules, creating ions in the process. e. Analyte molecules are heated with microwaves until they absorb sufficient energy to ionize.

14. Electrospray ionization is one technique used to interface chromatography with mass spectroscopy. Which statement regarding electrospray ionization is FALSE? a. The steel nebulizer is held at 0 V while the spray chamber is held at −3 500 V. b. The strong electric field between the nebulizer and spray chamber creates a fine aerosol. c. The glass capillary leading to the mass spectrometer is held at a potential of −4 500 V to attract the gas phase cations. d. Typically, the ions that vaporize from aerosol droplets were already charged in the solution from the chromatography column. e. There is minimal fragmentation, but fragmentation can be increased by collision-induced dissociation.

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Chap 22_10e 15. Which statements regarding the various mass spectrometry techniques are TRUE? I. Time-of-flight mass spectrometer – Ions expelled from the source all have the same kinetic energy.

For ions of different masses but equivalent kinetic energies, the lighter ions travel faster than heavier ions, separating the ions.

II. Quadrupole mass spectrometer – A constant voltage and a radio-frequency oscillating voltage are applied to the four metal rods. The electric field deflects ions in complex trajectories as the ions migrate from source to detector, allowing only one particular m/z to reach the detector.

III. Linear quadrupole ion-trap mass spectrometer – Sections are added at each end of the quadrupole to create a potential well. The ends are kept sufficiently positive relative to the center to trap cations in the center. The ions are confined in the radial direction by a radio-frequency field. By manipulating voltages, ions of specific m/z can be expelled.

IV. Orbitrap mass spectrometer – A central electrode is held at −5 kV and the outer electrodes are held near ground potential. A packet of ions is introduced parallel to the central electrode and an electric field pushes the electrons into orbit around the central electrode in the center of the orbitrap. a. I, II, and IV b. I, II, III, and IV c. II, III, and IV d. II and IV e. III and IV 16. When identifying the molecular ion peak M+•, all of the following guidelines must be kept in mind, EXCEPT: a. M+• will be the highest m/z value of any of the “significant” peaks in the spectrum that cannot be attributed to isotopes or background. b. intensities of isotopic peaks M + 1, M + 2, and so forth must be consistent with the proposed formula. c. M+• will be at least 70% of the base peak. d. the peak for the heaviest fragment ion should not correspond to an improbable mass loss from M+•. e. if a fragment ion is known to contain x atoms of element Z, then there must be at least x atoms of element Z in the molecular ion.

17. How are neutral species in solution converted to gaseous ions in atmospheric pressure chemical ionization? a. A high voltage is applied to a needle in the path of the aerosol. An electric corona forms around the charged needle, which injects electrons into the aerosol and creates ions. b. The neutral species is bombarded with 70-eV electrons. c. Neutral species are exposed to high-energy protons, which cause ionization by removing an electron from the neutral species. d. Only electrospray ionization can be used to convert neutral species into gaseous ions. e. It is not possible to convert a neutral species into gaseous ions. The species must already be charged in the solution from the chromatography column.

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Chap 22_10e 18. The molecular ion peak for a compound is at m/z 180. What number of nitrogen atoms are possible in the compound? a. an odd number of nitrogen atoms b. an even number of nitrogen atoms c. zero nitrogen atoms d. A and C e. B and C 19. In which area(s) of a mass spectrometer can collision-induced dissociation (collisionally activated dissociation) occur?

I. The entrance to the mass spectrometer II. The detector III. The reflectron IV. The collision cell V. The drift region a. IV b. III and V c. I and V d. I and IV e. III, IV, and V 20. Ion mobility spectrometry _______________ and __________________. a. operates at ambient pressure; is gas-phase electrophoresis b. operates at low pressure; is gas-phase electrophoresis c. operates at ambient pressure; is a form of mass spectrometry d. is gas-phase electrophoresis; is a form of mass spectrometry e. operates at high pressure; is a form of mass spectrometry 21. Electron ionization is one method that can be used to create ions from volatile analytes for analysis via mass spectrometry. Which statement describes how ions are created via electron ionization? a. Analyte molecules are bombarded with high-energy electrons (usually 70 eV), allowing the molecules to absorb enough energy to ionize. b. A reagent gas is converted to reactive species by energetic electrons. These new reactive species then donate a proton to the analyte molecules to produce the protonated molecule, MH+. c. Analyte molecules are ejected from the solution phase to the gas phase. During this process the analyte molecules lose an electron, causing ionization. d. Analyte molecules are bombarded with protons that pull electrons away from the analyte molecules, creating ions in the process. e. Analyte molecules are heated with microwaves until they absorb sufficient energy to ionize.

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Chap 22_10e 22. Calculate the number of rings plus double bonds for a compound with the chemical formula C23H28N2O5S2.

23. Electrospray time-of-flight mass spectrometry was used to analyze the eluate from a high-performance liquid chromatography separation. The mass spectrum of one chromatographic peak, containing a protein of unknown molecular mass, displays peaks at m/z = 5 654.208, 5 277.326, 4 947.590, 4 656.613, 4 397.990, and 4 166.576. Find the average molecular mass (M) of the neutral protein and its standard deviation.

24. Calculate the number of rings plus double bonds for a compound with the chemical formula C25H32N2O3.

25. An acceleration voltage of 25.0 kV is used to accelerate ions into the 1.24-m flight tube of a time-of-flight mass spectrometer. The frequency at which spectra can be recorded is dependent on how long it takes the slowest ion to travel from the source to the detector. If you wish to scan up to m/z 8 325, at what frequency (spectra/s) could spectra be recorded if a new spectrum begins each time the slowest ion reaches the detector?

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Chap 22_10e 26. The mass spectrum of an unknown compound has a molecular ion peak at m/z 194. The proposed chemical formula for the compound is C8H10N4O2. The M + 1 peak has a relative abundance of 10.1%. Calculate the theoretical intensity of the M + 1 peak. Does the actual intensity of the M + 1 peak support the proposed chemical formula?

Isotopic abundance Element factor for M + 1 H 0.012% C 1.08% N 0.369% O 0.038%

27. Calculate the resolving power for a mass spectrometer that successfully separates m/z 125.135 from 125.293.

28. Consider an element with two isotopes, x and y, whose natural abundances are a and b (a + b = 1), respectively. The probability of finding each combination of isotopes in a compound with two atoms of the element is derived from the expansion of the binomial (a + b)2, where the first term of the expansion corresponds to the probability of finding two atoms of isotope x in the compound, the second term of the expansion corresponds to the probability of finding one atom of each isotope in the compound, and the third term of the expansion corresponds to the probability of finding two atoms of isotope y in the compound.

Calculate the probability of finding each isotope combination of Br in CH2Br2. The natural abundance of 79Br is 50.69% and the natural abundance of 81Br is 49.31%.

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Chap 22_10e 29. A sample of copper containing two isotopes, 63Cu and 65Cu, was analyzed using a time-of-flight mass spectrometer. All of the ions were accelerated through a potential difference into a 0.800 0-m flight tube and have a kinetic energy of 1.000 × 10−16 J. The 63Cu+ ions took 1.829 × 10−5 s to reach the detector. How long will it take the 65Cu+ ions to reach the detector. The 65Cu+ ions have a mass of 1.079 × 10−25 kg.

30. What is the mass in daltons (Da) of an ion with a +1 charge that has a velocity of 6.945 5 × 104 m/s after it is accelerated through a potential difference of 20.0 kV?

31. Calculate the resolving power for m/z = 283.980 if the width at half-height is 0.076 m/z units.

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Chap 22_10e Answer Key 1. c 2. b 3. a 4. e 5. d 6. a 7. b 8. e 9. a 10. b 11. b 12. c 13. b 14. b 15. b 16. c 17. a 18. e 19. d 20. a 21. a 22. 11; Insert the number of carbon atoms, hydrogen atoms, and nitrogen atoms into c – h/2 + n/2 + 1.

23. 79 145.3 ± 0.4 Da; Use the equation Then use the equation

to determine the charge (n) associated with each peak. to determine the molecular mass associated with each peak. Average

the values and find the standard deviation.

24. 11; Insert the number of carbon atoms, hydrogen atoms, and nitrogen atoms into c – h/2 + n/2 + 1.

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Chap 22_10e 25. 19 400 spectra/s; Calculate the velocity of the ions using the equation

. Then use the length of the flight

tube and the velocity to calculate the drift time. One over the drift time is the number of spectra that can be collected per second.

26. The calculated intensity is 10.3%, which is within 10% of 10.1%, or 10.1% ± 1.0%. Multiply each isotopic abundance factor by the number of atoms for each element and take the sum.

27. 792; Insert the values into R = (m/z)/D(m/z).

28. The probability of two 79Br atoms is 0.256 9; the probability of one 79Br atom and one 81Br atom is 0.499 9; and the probability of two 81Br atoms is 0.243 1. The binomial expands to a2 + 2ab + b2. Insert the natural abundances in decimal form into the terms of the expanded binomial to calculate each probability.

29. 1.858 × 10−5 s; Solve the equation

for v. Divide the length of the flight tube by the velocity of the ions

to determine the drift time.

30. 807 Da; Rearrange the equation

to solve for m, which is the mass of the ion in kilograms. Then

convert the mass from kilograms to daltons.

31. 3 737; Insert the values into R = (m/z)/(m/z)1/2.

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Chap 23_10e Indicate the answer choice that best completes the statement or answers the question.

1. Which statements are NOT true regarding chromatography?

I. The stationary phase is typically a liquid bonded to the inside surface of a capillary or the surface of a solid packed in the column.

II. The mobile phase entering the column is the eluate.

III. The mobile phase is the solvent (gas or liquid) that moves through the column.

IV. Eluent is the process of passing mobile phase though a column.

V. Columns are either open tubular or packed. a. I and V b. II and IV c. I, II, and III d. III and V e. I, IV, and V 2. The constants A and C in the van Deemter equation barely contribute to band broadening in a. wall-coated open tubular columns. b. porous-layer open tubular columns. c. support-coated open tubular columns. d. packed columns. e. capillary electrophoresis.

3. Which statement(s) is/are TRUE for plate height?

I. The smaller the plate height, the narrower the peaks.

II. The larger the plate height, the better the separation between peaks.

III. Plate height can be calculated from the column length, L, and the number of theoretical plates, N. IV. For a fixed number of theoretical plates, N, the longer a solute is on the column, the wider the peak will become. a. I, II, and IV b. II and IV c. I, III, and IV d. II e. IV

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Chap 23_10e 4. Which form of chromatography is INCORRECTLY defined? a. adsorption chromatography – solute is adsorbed into solid particles b. partition chromatography – solute equilibrates between the stationary liquid and the mobile phase c. ion-exchange chromatography – solute ions of the opposite charge are attracted to the stationary phase d. size exclusion chromatography – solute is separated based on size e. affinity chromatography – solute binds to a molecule specific for the solute 5. Which statement regarding the diffusion coefficient is NOT true? a. The diffusion coefficient measures the rate at which molecules move randomly from a region of high concentration to a region of low concentration. b. The diffusion coefficient is related to the kinetic energy of a molecule and the friction it experiences while diffusing. c. Macromolecules diffuse 10 to 100 time more slowly than small molecules. d. A diffusion coefficient depends inversely on the viscosity of the medium. e. If elution time increases by a factor of 4, diffusion will broaden the band by a factor of 2.

6. The longer an analyte remains on a column, the broader the peak becomes. Peaks widen with time on column due to a. effusion. b. diffusion. c. adsorption. d. absorption. e. viscosity.

7. The van Deemter equation describes plate height in terms of the constants A, B, and C, and the linear velocity, ux. Which statement is NOT true for the van Deemter equation?

a. The constant A accounts for multiple pathways through the column. The value of A is column specific and independent of linear flow. b. The constant B accounts for longitudinal diffusion of the analyte in the mobile phase. The value of B is column specific, and the impact of B on the plate height is inversely proportional to the linear velocity. c. The constant C accounts for the finite time available for the solute to move between the stationary and mobile phases. The value of C is column specific, and the impact of C on plate height is proportional to the linear velocity. d. To minimize plate height, the optimal flow rate is the minima for the plot of B/ux + Cux versus flow rate. e. Linear velocity is the flow rate of the mobile phase.

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Chap 23_10e 8. Which statement regarding solute retention is NOT true? a. Retention time, t R, is the time that elapses between injection of the component onto the column and the arrival of the component at the detector minus the dwell time between the end of the column and the detector. b. Adjusted retention time is

, where t R is the retention time and t M is the time for the

mobile phase or unretained solutes to pass through the column. c. The separation factor, the sample, where d. The retention factor,

, is the ratio of the adjusted retention time for two components of so that α > 1. , is the time, t R, required to elute a component minus the time,

t M, for the mobile phase to pass through the column divided by t M. e. The separation factor, , is also equal to

, where KD2 and KD1 are the distribution constants

for components 2 and 1.

9. A solute is to be extracted using 150 mL of solvent. Which method will extract the maximum amount of solute? a. extract with 30 mL of solvent five times b. extract with 37.5 mL of solvent four times c. extract with 75 mL of solvent twice d. extract once with 150 mL of solvent e. extract with 50 mL of solvent three times 10. Which statement is NOT true when scaling up from analytical chromatography to preparative chromatography? a. Maintain a constant column length. b. The cross-sectional area of a column is inversely proportional to the mass of the analyte. c. Maintain a constant linear velocity. d. The sample volume applied to a column is proportional to the mass of the analyte. e. If you change the column length, the mass of sample can be increased in proportion to the total length.

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Chap 23_10e 11. A(n) ________________ is a plot of detector response versus retention time. a. spectrum b. electrophoretogram c. chromatogram d. chromospectrum e. elution response 12. The separation factor between two solutes, 1 and 2, is α = 1.35. If the adjusted retention time for solute 2 is 293 s, what is adjusted retention time for solute 1? a. 396 s b. 217 s c. 189 s d. 241 s e. 344 s 13. The retention factor for compound A on a column is 5.0. Given that t M = 1.0 min, what is the retention time for compound A? a. 5.0 min b. 6.0 min c. 4.0 min d. 1.0 min e. 0.2 min 14. What is the impact on the extraction of a weak acid into an organic solvent as the pH of the aqueous phase increases? a. The pH has no effect on the extraction of a weak acid into an organic solvent. b. As the pH increases, the value of D increases and the percentage of the weak acid left in the aqueous phase increases. c. As the pH increases, the value of D decreases and the percentage of the weak acid left in the aqueous phase decreases. d. As the pH increases, the value of D increases and the percentage of the weak acid left in the aqueous phase decreases. e. As the pH increases, the value of D decreases and the percentage of the weak acid left in the aqueous phase increases.

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Chap 23_10e 15. A Gaussian peak on a chromatogram is reported to have a width of 12.0 s. What is the peak width at halfheight? a. 3.0 s b. 6.0 s c. 7.1 s d. 5.1 s e. 0.78 s 16. Which comparison of open tubular columns to packed columns is INCORRECT? a. For similar analysis times, open tubular columns provide higher resolution than packed columns do. b. For the same length of column and applied pressure, the linear velocity in an open tubular column is higher than in a packed column. Therefore, the open tubular column can be made longer and still achieve a similar pressure drop and linear velocity. c. For similar analysis times, packed columns have decreased sensitivity compared to open tubular columns. d. Open tubular columns can achieve a resolution similar to packed columns in a shorter amount of time. e. Packed columns have a lower sample capacity than do open tubular columns.

17. The resolution between two peaks is defined as

Which variable is INCORRECTLY defined? a. t R is the retention time. b. VR is the retention volume. c. wav is the average peak width measured at the base of each peak. d. w1/2av is the average peak width measured at the base of each peak divided by 2. e. Resolution is the separation between two adjacent peaks.

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Chap 23_10e 18. The experimental data for the extraction of metal ions by dithizone into the ionic liquid 1-butyl-3methylimiddazolium hexafluorophosphate are given in the table.

pH 0.0 1.5 3.0 5.0 10.0

Percent extracted (%) Cu(II) Pb(II) 0 0 100 0 100 40 100 100 100 100

Cd(II) 0 0 5 22 100

What pH should be chosen to quantitatively extract Cu(II) from a mixture of Cu(II), Pb(II), and Cd(II)? a. pH 0.0 b. pH 1.5 c. pH 3.0 d. pH 4.5 e. pH 10.0 19. _______________ is the physical transfer of a solute from one phase to another. a. Partitioning b. Extraction c. Activated transport d. Sieving e. Miscibility 20. The observed variance,

, for a peak is the sum of the variances from all contributing factors. Which of

the following is NOT a contributing factor to

?

a. b. c. d. e.

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Chap 23_10e 21. The distribution constant for an amine, B, is KD = 3.0 and the acid dissociation constant for BH+ is Ka = 1.0 × 10−9. If 50.0 mL of a 0.010 M aqueous solution of the amine is extracted with 100.0 mL of solvent two times, what percentage of the amine remains in the aqueous solution if the pH of the aqueous phase is 9.00?

22. Calculate the number of theoretical plates and the plate height for a peak with a retention time of 1 295 s and a width at the baseline of 27 s that elutes from a 15 m column.

23. Separation of water and methanol was accomplished using a 1.0-m packed gas chromatography column.

An unretained species gave a sharp peak at 0.43 min, whereas water eluted at 1.28 min and methanol eluted at 1.53 min. Determine the linear velocity (in cm/s) of the mobile phase and the separation factor.

24. A chromatographic procedure was successfully developed to separate 5.0 mg of an unknown mixture on a column with a length of 60 cm and a diameter of 0.90 cm. What size of column should be used if 80 mg of the mixture are to be separated? If a 0.50 mL/min flow rate was used on the smaller column, what flow rate should be used on the larger column?

25. For two separated solutes, a = 1.10 and k 2 = 4.38. How many theoretical plates are required to achieve a resolution of 1.5?

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Chap 23_10e 26. A 0.10-mL sample was loaded onto an HPLC column. The detector volume was 0.15 mL. What are the variances due to injection and detection, respectively?

27. An unknown sample containing two components was separated on a column with a length of 15.0 cm.

Unfortunately, these two components were not well resolved, with a resolution of 1.2. To achieve a resolution of 2.0, what column length should be used? Assume that the separation factor and retention factors do not change greatly.

28. A compound eluted at 5.4 min from a column with a VM:VS ratio of 2:3, whereas an unretained species eluted at 1.1 min. Determine the distribution constant for the compound.

29. Calculate the weak acid (HA) concentration that remains in the aqueous phase when 50.0 mL of 0.184 M HA at pH = 4.75 is extracted with 50.0 mL of toluene three times. The distribution constant for the extraction is 5.17, favoring toluene. The Ka for HA is 8.0 × 10−4.

30. Calculate the resolution for two chromatographic peaks that elute at t R1 = 1 670 s and t R2 = 1 800 s with base peak widths of w1 = 51 s and w2 = 57 s.

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Chap 23_10e Answer Key 1. b 2. e 3. c 4. a 5. b 6. b 7. d 8. a 9. a 10. b 11. c 12. b 13. b 14. e 15. c 16. e 17. d 18. b 19. b 20. e 21. 6.25%; Calculate the distribution ratio, D, from [H+], Ka, and KD; substitute into the extraction equation; and solve for q2. Multiply q2 by 100 to get the percentage.

22. N = 3.7 × 104 plates; H = 0.41 mm; Calculate the number of plates, N, from the retention time and the baseline width using Equation 23-30. Then divide the number of plates into the length of the column to get the plate height.

23. ux = 3.9 cm/s; α = 1.29;

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Chap 23_10e 24. The newer column should be 60 cm in length and 3.6 cm in diameter; the flow rate should be 8.0 mL/min. The cross-sectional area of a column is proportional to the mass of analyte. The cross-sectional area of a column is proportional to the volume flow rate.

25. N = 6 600; Insert the values of a, k 2, and the resolution into the Purnell equation and solve for N. 26.

;

27. 42 cm; R ∝ N1/2, N = L/H; therefore, R ∝ L1/2 28. 2.6; First, determine k. Then use

to find KD.

29. 72.5% remains or 0.133 M HA remains in the aqueous phase; Calculate the distribution ratio, D, from [H+], Ka, and KD; substitute into the extraction equation; and solve for q3. Multiply q3 by the original concentration to calculate concentration remaining.

30. 2.41; Calculate Δt R and wav, and insert into Equation 23-23.

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Chap 24_10e Indicate the answer choice that best completes the statement or answers the question.

1. The temperature program for a separation starts at a temperature of 50o C and ramps the temperature up to 270o C at a rate of 10o C/min. Which statement regarding this separation is NOT true? a. At 10o C/min, a total of 22 min are needed to reach 270o C. b. Strongly retained solutes will remain at the head of the column when the temperature is low. c. Weakly retained solutes will separate and elute early in the separation. d. The vapor pressure of strongly retained solutes will increase as the temperature increases. e. Strongly retained analytes will give broad peaks.

2. Which carrier gas is most suitable for a gas chromatograph equipped with a thermal conductivity detector? a. H2 b. N2 c. CO d. CO2 e. air 3. Which statement regarding quantitative or qualitative analysis with gas chromatography is INCORRECT? a. The confirmation ion collected by a mass spectrometer is used to quantify a peak. b. Selected ion monitoring identifies a peak by monitoring the m/z characteristic for the compound. c. Internal standards quantify compounds by adding a known amount of a different compound to the sample and calculating the concentration from the peak areas for the compound and internal standard, and the response factor for the two compounds. d. Spiking identifies a peak by adding a known compound to the unknown. If the peak area increases, the peak has been identified. e. A linear response for concentration and detector response is required for quantitative analysis.

4. Cold trapping and solvent trapping are used with ________ injection to give _________ peaks. a. split; sharper b. on-column; faster eluting c. splitless; sharper d. split; faster eluting e. splitless; faster eluting

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Chap 24_10e 5. The first 3- to 10-m length of a capillary column serves as a guard column and retention gap. Guard columns a. improve peak shape by separating volatile solvent from less volatile solutes prior to separation on the chromatographic column. b. divert a portion of the sample to waste to improve peak shape. c. accumulate nonvolatile substances to prevent column contamination. d. begin the separation by retaining nonvolatile solutes but not volatile solutes. e. increase the number of theoretical plates, thus reducing the plate height.

6. Which statements regarding gas chromatography columns are TRUE? I. Most gas chromatography columns are open tubular columns composed of fused silica with a polyimide exterior coating.

II. Column inner diameters are typically 0.10 to 0.53 cm, and the columns are 15 to 100 cm in length, with a length of 30 cm being the most common.

III. A wall-coated open tubular column has a liquid stationary phase coating the inside wall of the column.

IV. A porous-layer column is filled with a highly porous stationary phase.

V. A packed column is a column filled with a solid stationary support coated in liquid stationary phase. a. I, III, and V b. II and IV c. III, IV, and V d. I and II e. I, II, III, IV, and V 7. Excessive column bleed should be avoided in gas chromatography. What are the causes of excessive column bleed?

I. Excessive column temperature II. Oxygen from leaking septum or fitting III. Oxygen contamination of carrier gas IV. Injection of chemicals that react with stationary phase a. I and III b. II an IV c. I and IV d. I, III, and IV e. I, II, III, and IV

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Chap 24_10e 8. The thermal conductivity detector functions by measuring the _____________ for a hot filament that is the result of a(n) _______________ to compensate for the ___________ of the filament by the mixture of analyte and carrier gas that elutes from the column. a. voltage increase; decreased resistance; warming b. voltage decrease; increased resistance; cooling c. voltage decrease; increased resistance; warming d. voltage increase; increased resistance; cooling e. voltage increase; increased resistance; warming 9. Which statement regarding gas chromatography carrier gases is NOT true? a. N2, He, and H2 are the most common carrier gases. b. Optimal linear velocity increases in the order N2 < He < H2. c. The fastest separations can be achieved with N2 as the carrier gas. d. N2 gives a lower detection limit than He with a flame ionization detector. e. Helium is the most common carrier gas and is compatible with most detectors.

10. Which statement regarding the techniques utilized in splitless injection for gas chromatographic analysis is NOT true? a. In solvent trapping, the initial column temperature is set 10°–20°C below the boiling point of the solvent. b. In cold trapping, the initial column temperature is at least 100°C below the boiling points of the solute of interest. c. In solvent trapping, both solvent and solutes are condensed at the column head. d. Cryogenic focusing is required, with an initial column temperature below room temperature, for low boiling. e. Solvent trapping and cold trapping sharpen the peaks for all components, including the solvent and solutes.

11. Compared to open tubular columns, packed columns provide a. sharper peaks. b. greater sample capacity. c. shorter retention times d. better resolution. e. all of the above.

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Chap 24_10e 12. As the temperature of a gas chromatography column increases, the ___________ of the analyte increases, the retention time ____________, and the peak shape ____________. a. vapor pressure; increases; sharpens b. solubility; decreases; sharpens c. vapor pressure; decreases; sharpens d. equilibrium; increases; broadens e. temperature; increases; sharpens 13. Which statement regarding flame ionization detectors is NOT true? a. The detector responds to hydrocarbon analytes, which are burned in a mixture of H2 and air after eluting from the column. b. During combustion the sample carbon atoms produce CH radicals, which are thought to produce CHO+ ions and electrons in flames. c. In the absence of analyte, a current of ~10−14 A flows between the flame tip and the collector.

Analytes produce a current of ~10−12 A. d. Flame ionization detectors respond to nonhydrocarbon compounds, but to a lesser degree than hydrocarbons. e. The detector responds proportionally to analyte mass over seven orders of magnitude.

14. Which statement regarding gas chromatography detectors is NOT true? a. Thermal conductivity detectors respond to the thermal conductivity difference between analyte and carrier gas. b. Flame ionization detectors are used only with inorganic compounds. c. Thermal conductivity detectors respond to all analytes. d. Flame ionization detectors generate signal by burning the analyte, which generates electrons. e. For both thermal conductivity and flame ionization detectors, detector response is proportional to analyte concentration.

15. Which statement regarding split injection is NOT true? a. The sample is completely vaporized and mixed in the mixing chamber by a combination of high temperature and brisk carrier flow. b. At the split point, a small fraction of vapor enters the column, but most passes through the waste vent. c. The proportion of sample that does not reach the column is the split ratio. d. Split ratios range from 1:1 to 500:1, and low split ratios give larger peaks. e. After injection, the waste vent is closed and the flow rate is increased.

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Chap 24_10e 16. Which of the statements regarding two-dimensional gas chromatography are TRUE? I. Two-dimensional gas chromatography is often called GC × GC. II. Two-dimensional gas chromatography uses two columns that have the same retention selectivity but different lengths to resolve component of complex samples.

III. The modulator positioned between the two columns accumulates eluate from the first column in a very short period of time and then rapidly admits the accumulated eluate onto the second column.

IV. The first column should be shorter than the second column so as to obtain a rapid separation in the second column.

V. For high resolution, the second column is narrower than the first column and has a thinner film of stationary phase than the first column. a. I, II, III, and V b. II, III, IV, and V c. II, V, III, I, IV d. I, III, and V e. I, III, IV, and V 17. Which statement regarding on-column injection for gas chromatographic analysis is NOT true? a. On-column injection is the best injection method for quantitative gas chromatographic analysis. b. There is no discrimination of analytes based on boiling point during the injection period. c. On-column injection should be limited to clean samples containing little or no nonvolatile components. d. A retention gap must be used for on-column injections of all samples. e. Programmed temperature vaporization is similar to on-column injection, and is used to facilitate the vaporization of volatile analytes from a liner.

18. Multiple decisions must be made when developing a new method using gas chromatography. Order the method development decisions in the order in which they should be made.

I. Detector II. Goal of analysis III. Injection IV. Sample preparation V. Column a. II, III, I, IV, V b. II, I, V, III, IV c. II, V, III, I, IV d. II, IV, I, V, III e. II, IV, III, I, V

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Chap 24_10e 19. Which technique is NOT a sample preparation technique used with gas chromatography? a. purge and trap b. solid-phase microextraction c. thermal desorption d. stir-bar sorptive extraction e. matrix matching 20. Which of the following is NOT a way to maximize analyte recovery when using solid-phase microextraction? a. select a fiber with a polarity similar to the analyte you are trying to recover b. agitate or stir the sample during the recovery c. use a thick or porous film for volatiles d. add 25% NaCl and adjust the pH, reducing the pH for acids and increasing the pH for bases e. use a large volume of headspace when analyzing the headspace above a sample 21. For quantitative analysis with gas chromatography a. the retention time reported by the instrument for a compound is proportional to the quantity of that compound. b. the area of a peak reported by the instrument for a compound is proportional to the quantity of that compound. c. the calculated response factor value for a compound is proportional to the quantity of that compound. d. the peak height reported by the instrument for a compound is inversely proportional to the quantity of that compound. e. the peak shape reported by the instrument for a compound is proportional to the quantity of that compound.

22. Under what conditions are split, splitless, and on-column injection used?

I. Split injection is preferred for samples where the analytes of interest constitute >0.1% of the sample.

II. On-column injection is used for samples that decompose above their boiling point.

III. On-column injection is preferred for qualitative analysis of gas chromatography samples.

IV. Split injection is preferred for samples with high-boiling-point solutes in low-boiling-point solvents.

V. Splitless injection is preferred for trace analysis of analytes that are less than parts-per-million in the sample. a. I, II, and III b. II, III, and IV c. II, IV, and V d. II, III, and V e. I, II, and V

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Chap 24_10e 23. _________________ columns are used with _______________ chromatography. a. Wall-coated open tubular; gas-solid b. Porous-layer open tubular; gas-solid c. Wall-coated open tubular; liquid-liquid d. Porous-layer open tubular; liquid-liquid e. Wall-coated open tubular; liquid-solid 24. The elution order for solutes from a gas chromatography column is loosely in order of a. decreasing molar mass. b. decreasing freezing point. c. decreasing dielectric constant. d. increasing boiling point. e. increasing entropy change.

25. Which statement regarding electron capture detectors is NOT true? a. The detector is sensitive to halogen-containing molecules, conjugated carbonyls, nitriles, nitro compounds, and organometallic compounds. b. The sensitivity of the detector is independent of the flow rate of the carrier gas. c. The detector uses 63Ni, a radioactive source, to emit high-energy electrons. d. Analytes capture some of the electrons and decrease the conductivity of the plasma. The detector responds by varying the frequency of voltage pulses to maintain a constant current. e. Each compound requires calibration for quantitation, as sensitivity varies widely for different analytes.

26. An unknown analyte was analyzed by gas chromatography along with n-heptane and n-octane. The retention times are given in the table. Calculate the Kovats retention index for the unknown.

Compound Retention time Methane (unretained) 1.32 min n-Heptane 4.65 min Unknown 7.21 min n-Octane 7.36 min

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Chap 24_10e 27. An alkane is used as an internal standard for the analysis of an alcohol. The response factor between the alcohol and the alkane is 2.54. The alkane is added to a sample of the alcohol with an unknown concentration. The alkane has a concentration of 25.9 nM in the sample and produces a peak area of 92 832.

The peak area for the alcohol in the sample is 23 422. What is the concentration of the alcohol in the sample?

28. To quantify the amount of ethanol in an organic sample, a technician prepares a standard solution containing 0.147 3 g of ethanol and 0.184 7 g of decane in 1.203 0 g of dichlormethane. Analysis by gas chromatography returns peak areas of 9 930 for ethanol and 7 893 for decane. Next, 0.9831 g of the unknown and 0.293 2 g of decane are added to 1.211 9 g of dichloromethane. Analysis by gas chromatography returns peak areas of 6 382 for ethanol and 11 283 for decane. What is the percent by mass ethanol in the sample?

29. A certain substance has an enthalpy of vaporization of 31.22 kJ/mol. At what Kelvin temperature will the vapor pressure be 5.00 times higher than at 309 K?

30. The enthalpy of vaporization of ethanol is 42.0 kJ/mol. The measured vapor pressure for ethanol at 18.0°C is 39.2 torr. Estimate the vapor pressure for ethanol at 62.0°C.

31. A column has a 0.217 mm diameter and a stationary phase thickness of 0.575 µm. An unretained solute passes through the column in 2.24 min, whereas the analyte requires 7.09 min. Calculate the distribution constant for the analyte.

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Chap 24_10e 32. A wall-coated open tubular chromatography column has a length of 30.0 m, an inner diameter of 250.0 mm, and a stationary phase thickness of 1.0 mm. Calculate the phase ratio of the column.

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Chap 24_10e Answer Key 1. e 2. a 3. a 4. c 5. c 6. a 7. e 8. e 9. c 10. e 11. b 12. c 13. d 14. b 15. e 16. d 17. d 18. d 19. e 20. e 21. b 22. e 23. b 24. d 25. b 26. I = 792; First, determine the adjusted retention times. Then insert the values into Equation 24-6 to find I. Copyright Macmillan Learning. Powered by Cognero.

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Chap 24_10e 27. 2.57 nM; The response factor times the ratio of the areas is equal to the ratio of the concentrations. Insert the values into the equation and solve for [X].

28. 10.69%; Use the mass of ethanol and decane in the standard solution, and the peak areas to calculate the response factor F. Next, use the response factor, the peak areas for ethanol and decane, and the mass of decane added to the unknown solution to calculate the mass of ethanol in the unknown. Divide the mass of ethanol by the mass of unknown and multiply by 100 to get the percent by mass ethanol.

29. 356 K; Use the following form of the Clausius-Clapeyron equation, where P2 = 5.00P1.

30. 383 torr; Use the Clausius-Clapeyron equation.

31. 204; First, determine the phase ratio and retention factor for the analyte. Then use these values to calculate the distribution constant.

32. 62.5; The phase ratio is defined as the mobile phase volume over the stationary phase volume. Use Equation 24-2 to calculate the phase ratio.

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Chap 25_10e Indicate the answer choice that best completes the statement or answers the question.

1. Which of the following are characteristics of reversed-phase chromatography?

I. The stationary phase is polar.

II. The mobile phase is more polar than the stationary phase.

III. A less polar mobile phase has a lower eluent strength.

IV. A more polar mobile phase has a higher eluent strength.

V. The stationary phase is nonpolar. a. I and IV b. II only c. I and III d. II and IV e. II and V 2. Which statement regarding a guard column is NOT true? a. The guard column contains the same stationary phase as the main column. b. The guard column should be replaced periodically. c. The guard column is used to remove fine particles. d. A guard column is highly cost effective for ≤ 5-cm columns. e. The guard column should be placed prior to the entrance to the main column.

3. Which statement regarding HPLC sample injection is NOT true? a. At least three loop volumes of sample should be flushed through the injector to ensure complete loop filling. b. The injection loop must be removed and cleaned after each injection. c. The sample solvent strength should be less than or equal to the eluent solvent strength. d. Samples should be passed through a 0.5-m filter to remove particulate matter prior to injection. e. A blunt syringe needle should be used to prevent damage to the injection port.

4. HPLC operation requires special, HPLC-grade solvents that require pretreatment before use. Which of the following is NOT a characteristic of HPLC-grade solvents or the pretreatment process required before use? a. The use of HPLC-grade solvents prevents the degradation of costly columns due to impurities. b. HPLC-grade solvents must be sparged prior to use if the instrument is not equipped with a degasser. c. HPLC-grade solvents minimize background contamination signals in the detector. d. HPLC-grade solvents must be filtered to remove particulate matter. e. HPLC-grade solvents must be stored in special cabinetry to prevent solvent degradation.

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Chap 25_10e 5. Which statement(s) regarding hydrophilic interaction chromatography (HILIC) is/are TRUE? I. HILIC is used to separate molecules that are too polar to be retained by reversed-phase columns.

II. HILIC uses a strongly polar stationary phase.

III. The elution order of analytes in HILIC is opposite to that in reversed-phase liquid chromatography. a. I b. II c. I and III d. I and II e. I, II, and III 6. Which statement regarding the effect of pH on the reversed-phase retention of weak acids and weak bases is NOT true? a. The ionized form is retained on the reversed-phase column, whereas the neutral form is weakly retained. b. The retention factor of a weak acid decreases as the pH increases. c. Retention is governed by the “average” form of the acid or the base. d. The retention factor of a weak base decreases as the pH decreases. e. The aqueous component of the mobile phase must contain a buffer.

7. Which detector is NOT used with HPLC? a. electrochemical detector b. ultraviolet detector c. refractive index detector d. evaporative light-scattering detector e. thermal conductivity detector 8. Which statement(s) is/are NOT true regarding partition and adsorption chromatography?

I. In adsorption chromatography, the lower the eluent strength of the solvent, the more easily it displaces the solute.

II. In partition chromatography, the more time the solute spends in the mobile phase, the easier it is to elute.

III. In partition chromatography, solute equilibrates between the mobile phase and the liquid phase bonded to the solid support.

IV. In adsorption chromatography, solvent molecules compete with solute molecules for binding sites on the stationary phase. a. IV b. I c. III and IV d. II e. I and II Copyright Macmillan Learning. Powered by Cognero.

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Chap 25_10e 9. When deciding between an isocratic elution or a gradient elution, a scouting gradient should be performed.

Which criterion based on the scouting gradient is INCORRECT? a. If Δt/t G < 0.25, use isocratic elution. b. If Δt/t G > 0.40, use gradient elution. c. For 0.25 < Δt/t G < 0.40, use either an isocratic or gradient elution. The decision depends on factors such as available equipment or the complexity of the sample. d. For Δt/t G < 0.25, a good starting isocratic elution solvent composition is the composition halfway through the Δt of the scouting gradient. e. For 0.25 < Δt/t G < 0.40, a good starting isocratic elution solvent composition is the composition halfway through the Δt of the scouting gradient.

10. Gradient elution can be used in HPLC to provide rapid elution of all components. How is the gradient created for a reversed-phase HPLC separation? a. The pH of the mobile phase is increased over time by adding a strong base. b. The pH of the mobile phase is decreased over time by adding a strong acid. c. The vapor pressure of the analytes is increased over time by increasing the temperature. d. The polarity of the mobile phase is increased over time by mixing with a more polar solvent. e. The polarity of the mobile phase is decreased over time by mixing with a less polar solvent.

11. A ______________ is a gradient developed to retain the resolution of a single gradient, but reduce the analysis time. a. broken gradient b. segmented gradient c. choppy gradient d. step gradient e. parcel gradient 12. Elution with a constant composition mobile phase is a. constant polarity elution. b. isopolar elution. c. isocratic elution. d. homogeneous elution. e. true elution.

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Chap 25_10e 13. The _________________ is a sensitive, almost universal detector with nearly equal response to equal masses of nonvolatile analytes. a. electrochemical detector b. charged aerosol detector c. ultraviolet detector d. evaporative light-scattering detector e. thermal conductivity detector 14. As particle size decreases in HPLC, the required solvent pressure ___________, the plate number _____________, and the resolution ______________. a. increases; increases; decreases b. decreases; decreases; increases c. increases; increases; increases d. increases; decreases; decreases e. decreases; increases; increases 15. Normal-phase chromatography uses a _________ stationary phase and a _________ mobile phase. The more __________ the mobile phase becomes, the stronger the eluent strength. a. polar; polar; polar b. nonpolar; nonpolar; polar c. polar; nonpolar; nonpolar d. polar; nonpolar; polar e. nonpolar; polar; nonpolar 16. Which of the following are attributes of a good separation for isocratic reversed-phase HPLC? I. The retention factors for all peaks between 0.5 and 20.

II. All peaks should be symmetric with an asymmetry factor B/A in the range 0.9–1.5.

III. For quantification, the minimum resolution between the two closest peaks should be 1.5, although 2.0 is best to compensate for small changes in separation conditions.

IV. The pressure must be less than or equal to 20 MPa to prolong the life of the instrument. a. I, II, III, and IV b. I, II, and IV c. I, II, and III d. II, III, and IV e. III and IV

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Chap 25_10e 17. A ________________ creates gradients from up to four solvents by proportioning liquids through a fourway valve before the pump, which pumps the mixture at high pressure into the column. a. high-pressure mixing pump b. multiple-source mixing pump c. low-pressure mixing pump d. gradient-mixing pump e. variable-composition mixing pump 18. The pressure required to push solvent through the column is dependent on several variables. Which relationship between column pressure and the corresponding variable is INCORRECT? a. Pressure is inversely proportional to column length. b. Pressure is directly proportional to solvent viscosity. c. Pressure is inversely proportional to particle diameter squared. d. Pressure is inversely proportional to column radius squared. e. Pressure is directly proportional to volume flow rate.

19. During method optimization, the separation factor can be adjusted by altering all of the following EXCEPT the a. column temperature. b. type of organic solvent. c. solvent composition. d. sample injection volume. e. column type.

20. In reversed-phase separations, ________ column volume(s) of initial solvent should be passed through the column after a gradient run to equilibrate the stationary phase with solvent for the next run. a. one b. two c. five d. ten e. twenty 21. Supercritical fluid chromatography provides _________ speed and ________ resolution, relative to liquid chromatography, because of _________ diffusion coefficients of solutes in supercritical fluids. a. increased; greater; higher b. increased; lower; higher c. decreased; greater; lower d. decreased; lower; lower e. increased; lower; lower

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Chap 25_10e 22. The _________________ responds to analytes that are significantly less volatile than the mobile phase. a. electrochemical detector b. ultraviolet detector c. refractive index detector d. evaporative light-scattering detector e. thermal conductivity detector 23. Which stationary phase is NOT a common bonded nonpolar stationary phase for HPLC? a. octadecyl b. phenyl c. pentafluorophenyl d. octyl e. amide 24. Calculate the polarity index of a mobile phase consisting of 10 vol% water and 90 vol% acetonitrile. The polarity index for water is 10.2 and the polarity index for acetonitrile is 5.8.

25. A 20-cm-long HPLC column is packed with 5.0-mm-diameter particles. The pressure required to have a 1.0 mL/min flow rate is 828 bar. What pressure is required to maintain the same flow rate if the packing particle size is decreased to 3.0 mm without changing other parameters?

26. A gradient elution was optimized on a 30- × 0.46-cm column using a flow rate of 0.75 mL/min. What flow rate should be used if the gradient conditions are transferred to a 15-× 0.39-cm column?

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Chap 25_10e 27. The dwell volume of an HPLC system used for a gradient separation is 2.5 mL. The flow rate is maintained at 1.00 mL/min and t M = 3.0 min. Calculate the time at which the start of the gradient would reach the detector.

28. A 30-cm-long HPLC column is packed with 4.6-mm-diameter particles. The pressure required to have a 1.0 mL/min flow rate is 980 bar. What pressure is required to maintain the same flow rate if the column is shortened to half the original length without changing other parameters?

29. What particle diameter is needed to give 15 000 plates over a 10.0-cm HPLC column?

30. What is the plate height for a 10.0-cm HPLC column packed with a 3.5-mm solid phase?

31. A wide gradient from 5 to 100% acetonitrile applied over 49.9 min is used to separate 12 compounds.

The sample is injected at t = dwell time. The first compound elutes from the column at 16.7 min, and the last compound elutes at 34.1 min. From this run, it was determined that gradient elution would be appropriate for the separation of this mixture of compounds. Based on these results, determine the appropriate starting and ending values for the percent acetonitrile in the solvent for the second run over the same gradient time.

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Chap 25_10e 32. A wide gradient from 5 to 100% acetonitrile applied over 40.2 min is used to separate 12 compounds.

The sample is injected at t = dwell time. The first compound elutes from the column at 13.5 min, and the last compound elutes at 27.5 min. Calculate Δt/t G. Based on the results of this first run, should isocratic or gradient elution be used to separate the mixture of compounds?

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Chap 25_10e Answer Key 1. e 2. d 3. b 4. e 5. e 6. a 7. e 8. b 9. e 10. e 11. b 12. c 13. b 14. c 15. d 16. a 17. c 18. a 19. d 20. d 21. a 22. d 23. e 24. 6.2; Use Equation 25-3 to calculate the polarity index of the mobile phase.

25. 2 300 bar; Equation 25-2 shows that the pressure is inversely proportional to the particle diameter, assuming all of the other parameters are held constant.

26. 0.27 mL/min; Calculate the volume of each column. Then use Equation 25-10 to calculate the flow rate.

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Chap 25_10e 27. 5.5 min; Calculate the dwell time and add it to t M. 28. 490 bar; Equation 25-2 shows that the pressure is proportional to the column length, assuming all of the other parameters are held constant.

29. 2.0 mm; Insert the values into Equation 25-1.

30. 1.17 × 10−3 cm or 11.7 mm; Combine Equation 25-1 with

to calculate the plate height.

31. from 36.7 to 69.9% acetonitrile; Use the following equation to calculate the initial and final percent ofacetonitrile.

32. Δt/t G = 0.348; gradient elution; Calculate Δt/t G. Because Δt/t G is less than 0.40, a gradient elution should be used.

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Chap 26_10e Indicate the answer choice that best completes the statement or answers the question.

1. A sample of NaNO3 and K2SO4 is separated on a suppressed-ion anion chromatography column. _____________ are retained and separated on the separator column. ______ is used to elute the analyte from the separator column. ____________ are removed in the suppressor to reduce the ________________ and allow for detection of analyte. a. Na+ and K+; HCl; and ; background absorbance b.

and

; HCl; Na+ and K+; background conductivity

c. Na+ and K+; KOH; d.

and

and

; background conductivity

; KOH; Na+ and K+; background absorbance

e. None of the above is correct.

2. A suppressor column in suppressed-ion chromatography a. converts the analyte to a nonionic form. b. converts the ionic eluent to a nonionic form. c. separates the analytes of interest from each other. d. increases the conductivity of the background eluent. e. decreases the conductivity of the ions of interest.

3. Which specific interactions can be utilized in affinity chromatography for biochemical applications? a. between antibodies and antigens b. between an enzyme and substrate c. between lectins and carbohydrates d. between ligands and receptors e. all of above 4. ___________________ is used to isolate a single compound or class of compounds from a complex mixture by passing the sample through a column where only one solute is bound. The remaining solutes pass through the column. The bound solute is eluted from the column by changing a condition, such as __________________, to weaken its binding. a. Adhesion chromatography; ratio of polar and nonpolar solvents b. Binding chromatography; the pH or ionic strength c. Affinity chromatography; the polarity of the solvent d. Adhesion chromatography; the temperature of the column e. Affinity chromatography; the pH or ionic strength

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Chap 26_10e 5. What is the normal order of elution in capillary zone electrophoresis? a. anions (highest mobility last), all neutrals (unseparated), cations (highest mobility first) b. anions (lowest mobility first), all neutrals (unseparated), cations (lowest mobility first) c. cations (highest mobility first), all neutrals (unseparated), anions (highest mobility last) d. cations (lowest mobility first), all neutrals (unseparated), anions (lowest mobility last) e. cations (highest mobility first), neutrals (lowest mobility first), anions (highest mobility first) 6. Which statements regarding size exclusion chromatography are TRUE? I. Size exclusion chromatography separates on the basis of size.

II. Small molecules freely pass through the stationary phase and have a Kav = 0.

III. Small molecules can fit in the pores of the stationary phase and effectively pass through a larger volume, so they elute last.

IV. The pores in the stationary phase are too small for large molecules to pass through, so large molecules elute first.

V. Size exclusion chromatography with a hydrophilic stationary phase and an aqueous solvent is called gel permeation chromatography. a. I, II, III, and V b. II and V c. I, II, and IV d. I, III, and IV e. III, IV, and V 7. Which is NOT true for the apparent, electroosmotic, and electrophoretic mobilities? a. μapp = μeo + μep b. c. d. e. uapp = (μeo + μep)E 8. Under what conditions will anions be detected at the cathode during capillary electrophoresis? a. μeo > μep b. μeo < μep c. μeo = μep d. μeo + μep > 1 e. μeo − μep < 1 Copyright Macmillan Learning. Powered by Cognero.

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Chap 26_10e 9. The detector commonly used with ion chromatography is the a. conductivity detector. b. fluorescence detector. c. ion detector. d. gravimetric detector. e. oxidation detector.

10. What type of exchange column should be used to separate a mixture of I−, F−, Cl−,

and

a. anion-exchange column b. cation-exchange column c. either a cation-exchange or an anion-exchange column d. neither a cation-exchange or an anion-exchange column e. More information is needed to determined which type of exchange column to use.

11. Capillary columns require conditioning prior to their first use. Typical conditioning is washing with 1 M NaOH for 1 h, followed by 1 h of water, then 1 h of 6 M HCl, and finally 1 h of run buffer. Which statements are INCORRECT? I. The NaOH wash is thought to generate the Si-OH groups on the silica surface.

II. The NaOH wash removes any acidic contaminants from the silica surface.

III. The HCl wash is thought to generate the Si-OH groups on the silica surface.

IV. The HCl wash removes metal ions from the capillary surface. a. I and III b. I and IV c. II and III d. II and IV e. I, II, and III 12. _________________ are instruments constructed on plastic or glass chips which use electroosmotic flow or pressure to move liquids through micrometer channels from one reaction site to another. a. Microosmotic devices b. Microcapillary devices c. Microelectrophoric devices d. Microfluidic devices e. Microscale devices

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Chap 26_10e 13. Which statements regarding electroosmotic velocity are TRUE? I. Electroosmotic velocity, ueo, is calculated experimentally using the migration time of a neutral molecule and the length of the column from the injector to the detector.

II. Electroosmotic velocity, ueo, is the product of the electroosmotic mobility, μeo, and the electric field, E. III. Electroosmotic velocity, ueo, is inversely proportional to electric field strength, E. IV. Electroosmotic velocity is equal to the electroosmotic flow. a. II and IV b. II, III, and IV c. I, III, and IV d. I and II e. I, II, and IV 14. The number of plates, N, in capillary electrophoresis is defined as

Which statement is

NOT true? a. As the apparent mobility, μapp, increases, the number of plates increases. b. As the total length of the capillary, Lt, increases, the number of plates increases. c. As the diffusion coefficient, D, increases, the number of plates decreases. d. As the distance to the detector, Ld , increases, the number of plates increases. e. As the voltage, V, decreases, the number of plates decreases.

15. _________________ occurs when an electric field is applied to a capillary that has an electric double layer at the negatively charged surface inside the capillary. a. Electroosmotic flow b. Electrophoretic flow c. Electrodynamic flow d. Electrodiffusion flow e. Electroeffusion flow

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Chap 26_10e 16. X−, Y−, and Z− are separated with an anion exchange column with a mobile phase containing A−. The strength with which the anions bind to the column is Y− > X− > Z− > A−. To elute the anions, the concentration of A− is ___________, with ________ eluting first, _______ eluting second, and ________ eluting last. a. decreased; Y−; X−; Z− b. increased; Z−; X−; Y− c. decreased; Z−; X−; Y− d. increased; Y−; X−; Z− e. increased; X−; Y−; Z− 17. Which statement regarding a lightly cross-linked ion-exchange column is TRUE? a. The lightly cross-linked column swells very little in water. b. The selectivity of the lightly cross-linked column is higher than that of a heavily cross-linked column. c. A lightly cross-linked column has a higher density of ion-exchange sites than a heavily cross-linked column. d. The lightly cross-linked column equilibrates faster than a heavily cross-linked column. e. The lightly cross-linked column has a higher exchange capacity compared to a heavily cross-linked column.

18. Which detector is NOT used with capillary electrophoresis? a. ultraviolet absorption b. conductivity c. electrospray mass spectrometry d. laser induced fluorescence e. evaporative light scattering 19. What impact does the pH of the mobile phase have on the ability of a strongly acidic exchanger to interact with ions? a. The ability of the exchanger to interact with ions decreases as pH decreases. b. The ability of the exchanger to interact with ions increases as pH increases. c. The pH has no impact on the ability of the exchanger to interact with ions. d. The ability of the exchanger to interact with ions decreases as pH increases. e. The ability of the exchanger to interact with ions increases as pH decreases.

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Chap 26_10e 20. Which statement regarding hydrophobic interaction chromatography is NOT true? a. Hydrophobic interaction chromatography has a hydrophobic stationary phase, which repels water and is not wetted by water. The technique is used primarily for protein purification. b. Proteins are salted out of solution with ammonium, sodium, and potassium salts of phosphate and sulfate. Thiocyanate, iodide, and perchlorate salts increase the solubility of proteins in water. c. The proteins are eluted from the column with a gradient of increasing salt concentration, increasing the solubility of proteins in water. d. The most common stationary phase is an agarose gel with hydrophobic phenyl or alkyl groups attached and a pore size large enough to accommodate proteins. e. A protein sample with a large concentration of ammonium sulfate is applied to the column. The ammonium sulfate causes the protein to stick to the hydrophobic surface of the stationary phase.

21. Which statement regarding capillary electrophoresis sample injection is INCORRECT? a. Hydrodynamic injection uses a pressure differential between the two ends of the capillary to inject a volume of sample into the capillary. b. Electrokinetic injection uses an electric field to inject a volume of sample into the capillary. c. Each analyte has a different app, so the composition of the loaded sample is different from the original sample when using electrokinetic injection. d. Hydrodynamic injection is most useful for capillary sieving electrophoresis, in which the liquid in the capillary is too viscous for electrokinetic injection. e. Stacking concentrates the cations and anions in the sample at the interface between the sample plug and run buffer.

22. Which of the following is NOT a pressure driven separation technique? a. ion-exchange chromatography b. affinity chromatography c. supercritical fluid chromatography d. capillary electrokinetic chromatography e. size exclusion chromatography 23. The retention time for an analyte separated via capillary electrophoresis is 115 s. The separation occurred with a 50-cm column, with the detector at 45 cm, and a potential of 30 kV. A neutral molecule took 200 s to reach the detector. What is the electrophoretic mobility for the analyte?

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Chap 26_10e 24. A sample was injected electrokinetically into a capillary column. The length of the sample plug inside the capillary was 2.5 cm, the injection electric field was 500 V/cm, the electric conductivity of the sample was 1/3 of that of the background electrolyte, and mapp = 4.0 × 10–8 m2/(V·s). How much time was required to perform the electrokinetic injection?

25. A 3.0-s hydrodynamic injection of a sample into a 53-cm capillary with a 75 mm diameter was performed using a 3.0 × 104-Pa pressure difference. Assume that the viscosity of the sample was 0.001 0 kg/(m·s).

What volume (in mL) of the sample was injected?

26. A size exclusion chromatography column has a diameter of 8.3 mm and a length of 29.0 cm. The volume between the particles is 43.0% of the total column volume. The solid portion of the particles occupies 14.0% of the total column volume, and the pores within the particles occupy the remaining 43.0% of the total column volume. Determine the volume at which molecules that are completely excluded would elute from the column and the volume at which the smallest molecules would elute from the column.

27. Na+ was separated from a mixture of cations via capillary electrophoresis, with negative detection. The separation occurred with a 50-cm column, with the detector at 45 cm and a potential of 30 kV. A neutral molecule took 200 s to reach the detector. The retention time for Na+ is 85 s. Calculate the number of theoretical plates for the Na+ peak. The diffusion coefficient for Na+ in water is 1.0 × 10−9 m2/s.

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Chap 26_10e 28. A Zorbacx® GF-250 size exclusion chromatography column with a pH 7.0, 0.2 M NaH2PO4 mobile phase was used to separate lysozyme and other compounds in a biological sample. The retention volume for lysozyme was 10.85 mL. Assume that the interstitial volume was 7.0 mL, and the total mobile phase volume in the column was 13.9 mL. What is Kav for lysozyme?

29. Capillary electrophoresis is used to separate two analytes. The retention times are 112 and 113 s on a 60cm capillary, with the detector 50 cm from the inlet. The voltage applied across the capillary is 20 kV. Assuming that the peaks eluted close enough that N is essentially constant, what is the resolution for the two peaks if N = 500 000 plates?

30. How much OH− is released when 25.00 mL of 12.0 mM

are quantitatively retained on an anion-

exchange column?

31. It took a neutral species 38 s to travel through the microfabricated channel in a polydimethylsiloxanebased microfluidic device under the influence of an electric field. Assume that the length of the channel was 4.2 cm, and the voltage applied was 1 200 V. What was the electroosmotic mobility?

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Chap 26_10e Answer Key 1. e 2. b 3. e 4. e 5. c 6. d 7. d 8. a 9. a 10. a 11. c 12. d 13. e 14. b 15. a 16. b 17. d 18. e 19. c 20. c 21. d 22. d

23. mep = 2.77 × 10−8 m2/Vs; Calculate uapp and divide by E, or calculate mapp with calculate meo with

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Next,

. mep = mapp − meo.

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Chap 26_10e 24. 4.2 s; The time is equal to the plug length divided by the speed,

The speed is calculated as

25. 0.13 mL; Calculate the volume using the following equation:

1 Pa = 1 kg/(m·s2)

26. 6.7 mL; 13 mL; Calculate the volume of the column. Determine the interstitial volume, which is equal to the volume at which the excluded molecules will elude, using the percent of the volume between particles. Determine the total volume of the mobile phase, which is equal to the volume at which the smallest molecules will elude.

27. 1.19 × 106 plates; Calculate mapp, then insert all values into

28. 0.56; Use the following equation to calculate Kav.

29. 1.57; First, determine the apparent speed for each analyte and then use the electric field to calculate the apparent mobility. Calculate the difference in mobilities, as well as the average mobility. Insert the mobility difference, average mobility, and number of plates into the resolution equation to solve for the resolution.

30. 0.600 meq; Each

to be retained releases two OH− from the column, and each mole of OH− has 1 equivalent

of charge. Equivalent charges = 2 (moles of

).

31. 3.9 × 10−8 m2/(V·s); A neutral species does not experience the electrophoretic flow; therefore,

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Chap 27_10e Indicate the answer choice that best completes the statement or answers the question.

1. A 100.00-mL solution of copper(II) chloride is reacted with excess silver nitrate to precipitate 8.926 0 g of silver chloride. Which calculation is NOT required to determine the molarity of the copper(II) solution? a. calculate the molar mass of silver chloride b. calculate the molar mass of copper(II) chloride c. calculate the moles of silver chloride from the precipitate mass d. calculate the moles of copper(II) chloride from the moles of silver chloride and the reaction stoichiometry e. divide the moles of copper(II) chloride by the volume of the copper(II) chloride solution expressed in liters 2. Combustion analysis of C, H, and S converts each element into a nonmetal oxide for analysis. The separation analysis of the CO2, H2O, and SO2 produced can be achieved using gas chromatography with a(n) I. UV absorbance detector.

II. infrared absorbance detector.

III. thermal conductivity detector.

IV. flame ionization detector. a. I and II b. II only c. III only d. II and III e. II, III, and IV 3. The final precipitation product should have a known, stable chemical composition. One way to obtain a known, stable chemical composition is via _____________, which is defined as the strong heating of a precipitate to change the chemical composition of the precipitate to a more stable form. a. drying b. ignition c. charring d. thermogravimetric analysis e. annealing 4. Which property is NOT a property of the ideal solid product for gravimetric analysis? a. insoluble b. known composition c. pure d. easily filterable e. high density Copyright Macmillan Learning. Powered by Cognero.

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Chap 27_10e 5. Crystallization occurs when precipitant is added to a solution, creating a temporary supersaturated solution.

Crystallization occurs in two stages: nucleation and particle growth. Which statement(s) is/are NOT true regarding nucleation and particle growth?

I. Nucleation occurs when solutes form clusters of sufficient size which reorganize into a more ordered structure capable of growing into larger particles.

II. Particle growth occurs when ions, molecules, or other nuclei condense onto a particle to form a larger crystal.

III. Sources of nucleation may be solution impurities or small scratches on a glass surface.

IV. Particle growth occurs when amorphous particles clump to create larger amorphous particles. a. I b. III and IV c. IV d. I, II, and IV e. II and IV 6. Homogeneous precipitation occurs when a precipitant is generated slowly from within an initially homogeneous solution by a chemical reaction. What is the main benefit of homogeneous precipitation? a. Nucleation dominates over particle growth when precipitation is slow to give smaller, purer, more easily filtered crystals. b. Particle growth dominates over nucleation when precipitation is slow to give smaller, purer, more easily filtered crystals. c. Nucleation dominates over particle growth when precipitation is slow to give larger, purer, more easily filtered crystals. d. Particle growth dominates over nucleation when precipitation is slow to give larger, purer, more easily filtered crystals. e. None of these statements describes the main benefit of homogeneous precipitation.

7. In __________________, a substance is heated, and its mass is measured as a function of temperature. a. mass spectrometry b. thermogravimetric analysis c. precipitation analysis d. combustion analysis e. electrogravimetric analysis

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Chap 27_10e 8. Ag+ (FM 107.868 2) is precipitated from solution using Na2CO3 (FM 105.988 8). Calculate the gravimetric factor to relate the mass of the product to the analyte. a. 0.782 b. 1.278 c. 0.391 d. 5.112 e. 0.196 9. Which statement regarding the precipitation of ionic compounds in the presence of an electrolyte is NOT true? a. Tiny colloidal particles coagulate into larger crystals in a solution containing excess electrolyte. b. Absorption occurs when the exposed ions on the colloidal surface absorb oppositely charged ions in the solution, making the particle surface charged. c. The charged particle surface attracts more oppositely charged ions and repels ions with the same polarity from the ionic atmosphere surrounding the particles. d. Colloidal particles must collide with one another to coalesce, but the ionic atmospheres of the particles repel one another. e. Heating promotes particle coalescence by increasing their kinetic energy to overcome electrostatic repulsion.

10. Combustion analysis of C, H, and S converts each element into a nonmetal oxide for analysis. Combustion analysis of O requires a different approach. Which step(s) is/are NOT required or are NOT accurate about the combustion analysis of O? I. Sample is thermally heated to decompose the compound in the absence of O2.

II. Gaseous products from pyrolysis are passed through nickelized carbon at 1 075°C to convert oxygen into CO. III. Gaseous products are sent to a gas chromatograph for analysis.

IV. Halogens (X) are converted to HX during pyrolysis, dissolved in water, and titrated with silver nitrate.

V. Gaseous products from pyrolysis are passed through nickelized carbon at 1 075°C to convert oxygen to CO2. a. III, IV, and V b. II and IV c. IV and V d. I, II, and III e. V

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Chap 27_10e 11. Which statement(s) is/are true regarding precipitate particle size?

I. Particle size must be large enough to not clog or pass through a filter.

II. Particle size is controlled by precipitation conditions.

III. Thermodynamics favors larger crystals because the surface atoms have higher energy than internal atoms, and the fraction of atoms on the surface is smaller for larger crystals than smaller crystals.

IV. Ostwald ripening is the process via which smaller particles redissolve and larger crystals grow. a. III b. I, II, and IV c. II and IV d. I, II, III, and IV e. II, III, and IV 12. Which of the following is NOT true for traditional combustion analysis of a hydrocarbon? a. During combustion all carbon is converted to carbon dioxide. b. During combustion all hydrogen is converted to water. c. The increase in mass for the trap containing P4O10 is proportional to the mass H in the compound. d. The increase in mass for the Ascarite II® trap is proportional to the mass C in the compound. e. During combustion all carbon is converted to carbon monoxide.

13. Modern combustion analysis uses dynamic flash combustion. What is the main reason dynamic flash combustion is employed? a. to limit the time oxygen flows freely through the analyzer b. to leave no residue in the ceramic crucible c. to guarantee the complete oxidation of select reaction products d. because the gas chromatograph requires the sample to be injected all at once e. to guarantee the complete reduction of select reaction products 14. _________________ is a process in which the sample is thermally decomposed in the absence of added O2. a. Dialysis b. Pyrolysis c. Digestion d. Decomposition e. Peptization

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Chap 27_10e 15. _________________ are solid particles that lack any crystalline order. a. Crystalline particles b. Amorphous particles c. Plastic particles d. Irregular particles e. Elastic particles 16. Precipitates often contain impurities and require treatment to remove the impurities. Which precipitation term is INCORRECTLY defined? a. gathering – precipitate collects a trace element of the solution b. masking agent – compound that prevents impurities from reacting with the precipitant c. inclusion – pocket of impurities trapped within the precipitate d. digestion – precipitates remain in mother liquor after precipitation to promote slow recrystallization of the product, form larger particles, and expel impurities e. coprecipitation – impurities are precipitated along with the desired product 17. In gravimetric analysis, the __________ of a product is used to calculate the quantity of the original analyte. a. concentration b. mass c. pressure d. density e. energy 18. Which of the following are techniques to promote particle growth?

I. Use a large volume of solution so that the concentrations of analyte and precipitant are low.

II. Rapidly add the precipitant to precipitate all analyte from solution. Then raise the temperature of the solution to promote crystal growth.

III. Slowly add precipitant with vigorous mixing to prevent a highly supersaturated condition in solution.

IV. Raise the temperature of the solution to increase solubility of the precipitate and decrease supersaturation. a. III and IV b. II, III, and IV c. II and IV d. I, II, and III e. I, III, and IV

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Chap 27_10e 19. The decomposition of ________ in boiling water is able to slowly produce OH− for homogeneous precipitation. a. urea b. thioacetamide c. dimethyl oxalate d. trimethyl phosphate e. sulfamic acid 20. ________________ are solid particles with long-range order. a. Crystalline particles b. Amorphous particles c. Rigid particles d. Regular particles e. Inelastic particles 21. The iron content of an iron supplement is to be determined. The manufacturer states that each tablet contains 1 000.0 mg of bioavailable iron (FM 55.845). A tablet with a mass of 4.928 3 g is digested in nitric acid to release Fe3+ into solution. After filtration, the iron in solution is precipitated as Fe(HCO3)3 (FM 238.895 5) using sodium bicarbonate (Na2HCO3, FM 84.007). What is the minimum mass of sodium bicarbonate that must be added to precipitate the iron in solution?

22. Combustion analysis of 0.193 8 g of a hydrocarbon increased the mass of the water trap by 0.275 0 g and the mass of the carbon dioxide trap by 0.597 3 g. What is the percent by mass for C and H?

23. A 4.89-g mixture containing only FeCl3 (FM 162.2) and AlCl3 (FM 133.34) was processed to convert the chlorides to the hydrous oxides, which were ignited to produce 2.02 g Fe2O3 (FM 159.69) and Al2O3 (FM 101.96). Calculate the mass percent of Fe (FM 55.845) and Al (FM 26.981) in the original mixture.

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Chap 27_10e 24. A solid mixture contains NaIO3 (FM 197.892 4) and Na2SO4 (FM 142.04). A 2.547 3-g sample of the mixture was dissolved in 100.0 mL of water. An excess amount of Ba(NO3)2 was added to precipitate both and . The precipitate weighed 3.762 6 g. Calculate the weight percent of both NaIO3 and Na2SO4 in the mixture.

25. A copper (FM 63.546) powder may have been partially oxidized to CuO (FM 79.545). A 1.284 2-g sample of the powder was heated in a stream of H2, and the final mass of the powder was 1.261 9 g.

Determine the purity of the copper sample.

26. A solution is prepared by dissolving 5.283 2 g of a solid mixture containing Ni (FM 58.693 4) and Cu (FM 63.546) in nitric acid to release Ni2+ and Cu2+. The nitric acid is neutralized and the solution is diluted to 200.0 mL. Next, excess sodium carbonate (FM 105.988 8) is added to the solution to form precipitate Ni2+ and Cu2+ as nickel(II) carbonate (FM 118.702) and copper(II) carbonate (FM 123.55). The precipitate has a total mass of 2.530 5 g. The precipitate is then ignited to form NiO (FM 74.692 8) and CuO (FM 79.545) with a total mass of 1.612 6 g. What is the percent by mass of Ni and Cu in the solid sample?

27. A 100.0-mL sample of a solution containing copper(II) (FM 63.546) is treated with excess sodium carbonate (FM 105.988 8) to precipitate 3.605 5 g of copper(II) carbonate (FM 123.55). What is the molarity of the solution?

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Chap 27_10e 28. Combustion analysis of 0.193 8 g of a hydrocarbon increased the mass of the water trap by 0.275 0 g and the mass of the carbon dioxide trap by 0.597 3 g. What is the empirical formula for the hydrocarbon?

29. How many milliliters of 3.00 M (NH4)2C2O4 are required to provide a 50% excess to completely precipitate Ca2+ from 100.0 mL of 0.234 M Ca2+?

30. The amount of Mg2+ in a solution can be determined by precipitation with phosphate and gravimetrically measuring the mass of Mg2P2O7 (FM 222.55) produced. If 0.385 g of Mg2P2O7 were produced from a 50.00-mL solution containing an unknown amount of Mg2+, what is the molarity of Mg2+ in the solution?

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Chap 27_10e Answer Key 1. b 2. d 3. b 4. e 5. c 6. d 7. b 8. a 9. b 10. c 11. d 12. e 13. d 14. b 15. b 16. c 17. b 18. e 19. a 20. a 21. 4.512 9 g NaHCO3; Calculate the moles of iron(III). Use the stoichiometry of the reaction to determine the moles of sodium bicarbonate needed to react with the amount of iron. Use the molar mass to calculate the mass.

22. 15.88% H and 84.12% C; Convert the masses of water and carbon dioxide to moles. Then convert moles of water to moles of H, and moles of carbon dioxide to moles of C. Next, convert moles of H to grams, and moles of C to grams. Divide the mass of each element by the mass of the compound to get the percentage.

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Chap 27_10e 23. Fe, 9.4%; Al, 14.7%; Set up two equations, where x is the mass of Fe and y is the mass of Al.

Solve for x and y to get the mass of each element. Then divide by the sample mass and multiply by 100 to determine the mass percent.

24. NaIO3, 39.81%; Na2SO4, 60.19%; Set up two equations

Solve for x and y to get the mass of each salt. Then divide by the sample mass and multiply by 100 to determine the weight percent.

25. 91.37%; The change in mass of the powder is due to the loss of O in CuO. Therefore, the mass loss is used to determine the amount of CuO in the sample. The purity is based on the mass percent of Cu. 26. 11.14% Ni, 13.07% Cu; Set up two equations, x + y = 1.612 6 g for the metal oxides and (118.702/74.692 8)x + (123.556/79.545)y = 2.530 5. Solve for x and y to get the mass of each metal oxide. Convert the mass of each metal oxide to the mass of each metal. Divide by the sample mass and multiply by 100 to determine the mass percent.

27. 0.291 8 M Cu2+; Determine the moles of copper(II) carbonate from the mass and formula mass. Use the reaction stoichiometry to determine the moles of copper(II) in the solution. Divide by the volume of the solution to determine the molarity.

28. C4H9; Convert the masses of water and carbon dioxide to moles. Then convert moles of water to moles of H, and moles of carbon dioxide to moles of C. Divide by the smallest number of moles. Then multiply the 2.25 on H and the 1 on C by 4 to get C4H9.

29. 11.7 mL; Determine the moles of Ca2+ from the volume and concentration of the solution. Use the stoichiometry of the reaction to determine the moles of required to completely react with Ca2+. Then find the total volume of oxalate (required plus 50% excess).

30. 0.069 2 M; Determine the moles of Mg contained in the precipitate. Then divide the moles of Mg by the volume of the sample.

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Chap 28_10e Indicate the answer choice that best completes the statement or answers the question.

1. Which statement(s) is/are true regarding sample storage?

I. Steel needles should be avoided in biochemical analysis to prevent metal contamination.

II. Glass storage bottles may alter the concentrations of trace ions and proteins in solution through ion exchange.

III. Plastic storage bottles may absorb trace levels of analyte from solution. a. I b. II c. III d. I and II e. I, II, and III 2. Which statement regarding the dissolution of inorganic materials in acid is NOT true? a. Metals with positive reduction potentials should dissolve in nonoxidizing acids, such as HCl, HBr, HF, and H3PO4. b. Volatile species formed by protonation of anions, such as carbonate, will be lost from hot acids in open vessels. c. Volatile metal halides, such as SnCl4, and some metal oxides can be lost from hot acids in open vessels. d. Hot hydrofluoric acid dissolves silicates. e. Substances that do not dissolve in nonoxidizing acids may dissolve in oxidizing acids such as HNO3, hot concentrated H2SO4, or hot concentrated HClO4.

3. ____________ is the procedure in which an analyte is chemically modified to make the compound easier to detect or separate. a. Labeling b. Derivatization c. Tagging d. Marking e. Modification 4. Aqua regia solution, which consists of _______________, is able to dissolve Au and Pt. a. 3:1 (vol/vol) of concentrated HCl:HNO3 b. 3:1 (vol/vol) of concentrated HNO3:HCl c. 3:1 (vol/vol) of concentrated HCl:H2SO4 d. 3:1 (vol/vol) of concentrated H2SO4:HCl e. 1:1 (vol/vol) of concentrated H2O2:H2SO4 Copyright Macmillan Learning. Powered by Cognero.

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Chap 28_10e 5. Which techniques are NOT liquid extraction techniques?

I. Emulsive liquid-liquid microextraction II. Dispersive liquid-liquid microextraction III. Solid-supported liquid-liquid extraction IV. Supercritical fluid extraction V. Solid-liquid microextraction a. I and II b. III and V c. I and IV d. IV only e. I and III 6. To calculate the number of required replicate analyses, the equation

is used. Which statement

regarding the equation and calculation is NOT true? a. t is the Student’s t value for n – 1 degrees of freedom. b. ss is the sampling standard deviation. c. e is the difference between the sampling standard deviation and the desired sampling standard deviation. d. The initial calculation assumes an infinite n value and uses the corresponding Student’s t value. e. The n value from a calculation dictates the Student’s t value used in the next calculation.

7. Which of the following is NOT used to grind solids into a fine powder? a. mortar and pestle b. Wig-L-Bug® c. Shatterbox laboratory mill d. ball mill e. vortexer 8. QuEChERS is a sample preparation procedure involving extraction and sample cleanup. What does QuEChERS stand for? a. Quite, Easy, Cheap, Effective, Rugged, and Safe b. Quiet, Easy, Chelating, Effect, Routine, and Save c. Quite, Effective, Chelating, Effect, Running, and Safe d. Quick, Easy, Cheap, Effective, Rugged, and Safe e. Quiet, Efficient, Chopping, Effect, Runs, and Save

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Chap 28_10e 9. You can prepare solid samples of a certain particle size by grinding larger particles and passing them through a mesh sieve. Which solid sample has the smallest sampling variance if the mass of each sample is the same? a. 6/8 mesh b. 10/12 mesh c. 16/20 mesh d. 60/80 mesh e. 120/200 mesh 10. _______________ uses a hot, molten inorganic _______________ to dissolve substances that will not dissolve in acid. a. Smelting; solvent b. Fusion; flux c. Smelting; flux d. Fusion; solvent e. None of these choices is correct.

11. Which statement(s) is/are true regarding the statistics of sampling?

I. The total standard deviation is the sum of the analytical standard deviation and the sampling standard deviation.

II. If either sa or ss is sufficiently smaller than the other, there is little point in trying to reduce the smaller one.

III. Sample uncertainty may be reduced with a larger test portion. a. I b. II c. III d. II and III e. I, II, and III 12. Sample preparation may include a. removing interferents from the sample. b. derivatizing the analyte into a form suitable for easy measurement. c. concentrating the analytes. d. converting a bulk sample into a homogeneous laboratory sample. e. all of above.

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Chap 28_10e 13. Which sampling term is incorrectly defined? a. laboratory sample – smaller, homogeneous sample taken from the bulk; has the same composition as the bulk b. lot – the total material from which samples are taken c. aliquot – small portion of the bulk sample taken for individual analysis d. sample preparation – the series of steps that convert a representative bulk sample into a form suitable for analysis e. bulk sample – representative sample taken from the lot for analysis or archiving 14. Which statement regarding the decomposition of organic substances is NOT true? a. Microwave-induced combustion is a form of dry ashing. b. An example of wet ashing is microwave digestion with acid in a Teflon bomb. c. Wet ashing with perchloric acid is an extremely hazardous procedure and is best avoided. d. Kjeldahl digestion digests organic material with sulfuric acid for nitrogen analysis. e. The Carius method of wet ashing is performed with fuming H2SO4 in a sealed, heavy-walled glass tube at 200° to 300°C. 15. For random errors, the overall variance, variance,

, is calculated from the analytical variance,

, and the sampling

. What is the equation for calculating the overall variance?

a. b. c. d. e.

16. A mixture contains 5% NaCl particles and 95% NaBr particles. If 103 particles are taken, what is the relative standard deviation for NaCl and NaBr? a. 13.78% for NaCl; 0.73% for NaBr b. 186.2% for NaCl; 9.80% for NaBr c. 95.00% for NaCl; 5.00% for NaBr d. 54.39% for NaCl; 2.86% for NaBr e. 6.89% for NaCl; 0.36% for NaBr

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Chap 28_10e 17. Sample storage is an important aspect of sample preparation. Which statement regarding sample storage is NOT true? a. Glass containers should not be used to store samples for trace analysis of ions and proteins in solution. b. Analytes may react with the storage container. c. A stored sample may change with time. d. New plastic containers do not require washing before use. e. A stored sample may react with air.

18. Which statement regarding the preparation of a laboratory sample for analysis is INCORECT? a. Laboratory samples that do not dissolve under mild conditions may require acid digestion or fusion. b. Organic material may be destroyed by combustion (dry ashing) or wet ashing. c. Solids are typically dried at 110°C at atmospheric pressure to remove absorbed water prior to analysis. d. A portion of a coarse solid laboratory sample should be ground into a fine powder for analysis. e. Temperature-sensitive samples may be stored in an environment to achieve a constant, reproducible moisture level.

19. __________________ uses a small volume of chromatographic stationary phase or molecularly imprinted polymer to isolate analytes from the sample matrix to simplify analysis. a. Solid-phase extraction b. Polymer extraction c. Chromatographic extraction d. Antigen extraction e. Solvent-free extraction 20. ___________________ utilizes a surfactant to extract and preconcentrate analytes from a liquid sample. a. Dispersive microextraction b. Solid-supported liquid-liquid extraction c. Cloud point extraction d. Solid-phase microextraction e. Supercritical fluid extraction 21. On average, talcum powder contains approximately 32% MgO. How many particles of 325/1 250 mesh (average diameter 25 µm) sample are required to reduce the MgO sampling uncertainty to 1.0%?

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Chap 28_10e 22. What mass of a solid sample containing 2.5% analyte particles and 97.5% inert particles must be taken to give a standard deviation of 2% of the number of analyte particles? The average particle volume is 2.54 nL and the sample has an average density of 1.548 g/mL.

23. A mixture contains 2.5% KBr and 97.5% inert particles. What is the relative standard deviation for KBr if 50 000 particles are taken?

24. A 0.85-g sample has a 4% sampling standard deviation. What mass of sample must be taken to decrease the sampling standard deviation to 1%?

25. Calculate the mass corresponding to 1 500 spherical particles of a mixture that is 82% Na2CO3 and 18% NaHCO3. The average diameter of the particles is 0.85 mm. The density of Na2CO3 is 2.532 g/mL and the density of NaHCO3 is 2.159 g/mL.

26. A sample made of 1-mm-diameter particles contains 50 000 inert particles and 600 KNO3 particles. What is the expected number of KNO3 particles if a random sample of 5 000 particles is taken from the sample?

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Chap 28_10e 27. How many replicate analyses must be completed for a sampling deviation of 3.5% to achieve a 1% standard deviation?

28. A 5.8-g sample of liver homogenate was needed to give a standard deviation of ±2.4%. What is the sampling constant?

29. What mass will give a sampling standard deviation of 3.5% for a sampling constant of 40.9?

30. Calculate the overall standard deviation if the analytical standard deviation is 6% and the sampling standard deviation is 4%.

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Chap 28_10e Answer Key 1. e 2. a 3. b 4. a 5. c 6. c 7. e 8. d 9. e 10. b 11. d 12. e 13. c 14. e 15. e 16. a 17. d 18. d 19. a 20. c 21. 2.1 × 104 particles; Solve

for n, where p = 0.32 and q = 0.68.

22. 0.383 4 g; First, solve for n using the equation

, where p = 0.025 and q = 0.975, 0.000 4pn2 –

pqn = 0. The equation is a quadratic with c = 0, (y = mx + b). The value of n = 97 500. Multiply by the average volume per particle, then convert nanoliters to liters and liters to milliliters. Multiply by the density to get the mass in grams.

23. 2.79%; Calculate the number of KBr particles expected to be in 50 000 particles (1 250 KBr particles). The standard deviation is 34.91. Divide the standard deviation by 1 250 and convert to a percentage.

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Chap 28_10e 24. 13.6 g; First, calculate Ks using the equation Ks = mR2. Then calculate the mass, m, from Ks /R2.

25. 1.2 g; First, calculate the average density. Then calculate the mass of one particle and the total mass of all particles.

26. 59 KNO3 particles; Expected number of KNO3 particles 27. 50; Assume n = infinite and use Student’s t = 1.960, s = 0.035, and e = 0.01 to obtain n = 47. Then repeat the calculation using the Student’s t value of n − 1 = 46, to calculate n = 50. Repeating the calculation for 49 degrees of freedom yields the same value n.

28. 33; Use the equation Ks = mR2.

29. 3.34 g; Calculate the mass, m, from Ks /R2.

30. 7.2%; Solve the equation

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for so .

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