Introductory Chemistry Version 2.0 2nd Edition By David W. Ball TEST BANK

Chapter 1 What Is Chemistry?

True/False 1. Chemistry is the study of the interactions of matter with other matter and with energy.

True; Easy 2. Matter is anything that has a mass and takes up space.

True; Easy

3. Air found inside an empty chamber cannot be classified as matter.

False; Easy 4. Physical properties are characteristics of matter that describe how matter changes form in the presence of other matter.

False; Easy 5. Chemical properties are characteristics that describe matter as it exists.

False; Easy 6. A physical change occurs when a sample of matter changes one or more of its physical properties.

True; Easy 7. A burning sparkler is an example of chemical change.

True; Easy 8. Ice melting is an example of a chemical change.

False; Easy 9. A sample of matter that has the different physical and chemical properties throughout is called a substance.

False, Easy 10. An element can be broken down into simpler chemical substances by chemical reactions.

False; Easy 11. A compound is a combination of more than one element.

True; Easy 12. A compound‘s physical and chemical properties are different than the physical and chemical properties of its constituent elements.

True: Easy 13. Table salt has the same physical properties as its constituent elements, sodium and chlorine gas.

False; Medium

14. Carbonated water, before it is opened, is an example of heterogeneous mixtures.

False; Easy (the CO2 dissolved in the water makes it homogeneous. However, students will think of the water after it is opened, the CO2 bubbles are present making it look like it is heterogeneous.

15. Solution is another word used to describe homogenous mixtures.

True; Easy 16. A metal is an element that is brittle at room temperature.

False; Easy 17. Elements that have properties of both metals and nonmetals and are called antimetals.

False; Easy 18. Nonmetals do not conduct electricity or heat very well.

True; Easy 19. Science is the process of knowing about the natural universe through observation and experiment.

True; Easy 20. A hypothesis is a general statement that explains a large number of observations.

False; Easy 21. Experiments are not performed if a hypothesis exists on the issue to be tested.

False; Easy 22. A law is a specific statement that is thought to be never violated by the entire natural universe.

True; Easy 23. Science is concerned only with the natural universe.

True; Easy 24. Physics and astronomy are scientific fields concerned with the fundamental interactions between matter and energy.

True; Easy 25. Chemistry is widely regarded as the language of science.

False; Easy 26. A qualitative description implies a description of the extent to which a compound is used in a reaction.

False; Easy 27. The temperature outside is 95 degrees Fahrenheit; this is an example of a quantitative description.

True; Easy (students might not know the color of rhodochrosite unless they look it up)

28. A qualitative description means knowing how much of something is present.

False; Easy Multiple Choice Questions 29. Which of the following statements is the correct description of chemistry? a. It is the study of the interactions of matter with other matter and with energy. b. It is the study of the static behavior of particles and substances. c. It is the study of conversion of one form of energy to another form. d. It is the science that deals with kinetic and potential energy. e. It is the science that deals with matter, energy, and force. a; Easy 30. Anything that has mass and takes up space is known as _____. a. gravity b. matter c. leverage d. vacuum e. force b; Easy 31. Which of the following is not classified as matter? a. Air in a room b. Gas in a chamber c. Blood in blood vessels d. Thoughts of a human brain e. Air in human lungs d; Easy 32. Which of the following statements describes a chemical property? a. Mercury is the only metal that exists as a liquid at room temperature. b. Women usually have long nails. c. The stereo system is silver in color. d. A kilogram of iron weighs one thousand grams. e. Sodium reacts with water to produce sodium hydroxide and hydrogen. e; Moderate 33. Chemical properties are characteristics of matter that describe _____. a. how matter changes form in the presence of other matter b. how matter changes form when it releases kinetic or static energy c. how matter changes form when it acquires static energy d. how matter changes form when it acquires kinetic energy e. how matter changes form when it is heated or cooled a; Easy 34. Which of the following is an example of chemical change? a. Ice melts to form water. b. Water is turned to steam by heating. c. Temperature of the engine goes up by 5oC. d. Alkanes have less mass than alkenes.

e. Alkanes burn in the presence of oxygen. e; Easy 35. A burning match stick is an example of a(n) _____ change. a. physical b. mechanical c. chemical d. ionic e. nuclear c; Easy 36. A sample of matter that has the same physical and chemical properties throughout is called a _____. a. form b. base c. substance d. particle e. structure c; Easy 37. A(n) _____ is the simplest type of chemical substance; it cannot be broken down into simpler chemical substances by ordinary chemical means. a. compound b. element c. mixture d. base e. structure b; Easy 38. A _____ is a combination of more than one element. a. physical change b. phase c. thought d. chemical change e. compound e; Easy 39. Which of the following is an example of a heterogeneous mixture? a. Oxygen dissolved in water b. Carbon dioxide dissolved in water c. Sodium chloride in water d. A combination of salt and steel wool e. A combination of oxygen and hydrogen d; Moderate 40. A(n) _____ is an element that is solid at room temperature, is shiny and silvery, conducts electricity and heat well, can be pounded into thin sheets, and can be drawn into thin wires. a. sub-atom b. metal c. ion

d. semimetal e. nonmetal b; Easy 41. Which of the following elements‘ property allows them to be pounded into thin sheets? a. Conductivity b. Resistivity c. Malleability d. Ductility e. Impeditivity c; Easy 42. Which of the following elements‘ property allows them to be drawn into thin wires? a. Malleability b. Ductility c. Conductivity d. Resistivity e. Impeditivity b; Easy 43. A nonmetal is an element that is _____. a. brittle when solid b. conductive to electricity c. hard when solid d. conductive to heat e. characterized by high ductility a; Moderate 44. Which of the following terms refers to the elements that have properties of both metals and nonmetals?

a. Base metals b. Precious metals c. Noble metals d. Metalloids/Semimetals e. Extracted metals d; Easy 45. _____ can be regarded as the process of knowing about the natural universe through observation and experiment.

a. Science b. Civics c. History d. Psychology e. Geography a; Easy 46. Which of the following terms refers to an educated guess about how the natural universe works? a. Experimentation b. Empirical analysis

c. Hypothesis d. Postdiction e. Intuition c; Easy 47. Which of the following is the first step in a scientific method? a. Forming a theory b. Stating a hypothesis c. Refining a hypothesis d. Validating assumptions e. Validating a theory b; Easy 48. Which of the following refer to tests of the natural universe to see if a hypothesis is correct? a. Sequences b. Topology c. Internship d. Amendments e. Experiments e; Easy 49. A(n) _____ is a general statement that explains a large number of observations. a. hypothesis b. model c. purpose d. theory e. experiment d; Easy 50. A specific statement that is thought to be never violated by the entire natural universe is called a(n) ______. a. premise b. hypothesis c. law d. theory e. observation c; Easy 51. The fact that all matter attracts all other matter is an example of a(n) _____. a. law b. observation c. argument d. purpose e. hypothesis a; Easy 52. Which of the following scientific fields is known as the language of science? a. Physics b. Mathematics c. Chemistry

d. Botany e. Zoology b; Easy 53. A _____ description implies a description of the features of an object. a. static b. quantitative c. qualitative d. descriptive e. dynamic c; Easy 54. Which of the following types of descriptions represents the specific amount of something? a. Qualitative b. Descriptive c. Static d. Dynamic e. Quantitative e; Easy 55. The compound sulfur is yellow colored. This description of sulfur is an example of a _____ description. a. Qualitative b. Dynamic c. Static d. Quantitative e. Descriptive a; Easy Essay Questions 56. What is matter? Explain by providing examples.

Matter is anything that has mass and takes up space. A book, a computer, and food are examples of matter. Air is also an example of matter because it occupies space.

Easy 57. What is meant by physical properties of matter?

Physical properties are characteristics that describe matter as it exists. Some of many physical characteristics of matter are shape, color, size, and temperature. An important physical property is the phase (or state) of matter. The three fundamental phases of matter are solid, liquid, and gas.

Easy 58. What is meant by chemical properties of matter?

Chemical properties are characteristics of matter that describe how matter changes form in the presence of other matter. Burning is an example of a chemical property.

Easy 59. What is physical change? Provide an example.

A physical change occurs when a sample of matter changes one or more of its physical properties. For example, a solid may melt, or alcohol in a thermometer may change volume as the temperature changes. A physical change does not affect the chemical composition of matter.

Easy 60. From the following, which is not an example of a chemical change? Be sure to explain your reasoning.

1. Burning a plastic water bottle 2. Production of hydrogen gas from water 3. Tarnishing a copper penny 4. Chopping a log into sawdust 5. Charging a cell phone A chemical change is the process of demonstrating a chemical property. Chopping a log into sawdust is an example of a physical change and not a chemical change. The log is being chopped into smaller pieces, but is not changing the chemical composition of the log. The other choices are all examples of chemical changes. Chemical changes are frequently accompanied by physical changes, as the new matter will likely have different physical properties from the original matter.

Moderate 61. What is called a substance? Provide an example.

A sample of matter that has the same physical and chemical properties throughout is called a substance. Examples are carbon blocks, water etc.

Easy 62. What are elements? Compare elements with compounds.

An element is the simplest type of chemical substance. It cannot be broken down into simpler chemical substances by ordinary chemical means. Each element has its own unique set of physical and chemical properties. Examples of elements include iron, carbon, and gold. A compound is a combination of more than one element. The physical and chemical properties of a compound are different from the physical and chemical properties of its constituent elements.

Moderate 63. What are mixtures? Write the types of mixtures, and then give an example for each type.

Physical combinations of more than one substance are called mixtures. There are two types of mixtures. A heterogeneous mixture is a mixture composed of two or more substances (oil and water). It is easy to tell, sometimes by the naked eye, that more than one substance is present. A homogeneous mixture is a combination of two or more substances that is so intimately mixed that the mixture behaves as a single substance (salt water).

Moderate 64. What is a metal? How is it different from a nonmetal?

A metal is an element that is solid at room temperature, is shiny and silvery, conducts electricity and heat well, can be pounded into thin sheets, and can be drawn into thin wires. A nonmetal is an element that is brittle when solid, does not conduct electricity or heat very well, and cannot be made into thin sheets or wires. Nonmetals also exist in a variety of phases and colors at room temperature.

Easy 65. Explain the concepts of malleability and ductility.

Metals can be pounded into thin sheets. This property is called malleability.

Metals can be drawn into thin wires. This property is called ductility.

Easy 66. What is science? What is its importance?

Science is the process of knowing about the natural universe through observation and experiment. Science is not the only process of knowing, but it has evolved over more than 350 years into the best process that humanity has devised to date to learn about the universe around us.

Easy 67. What is a hypothesis? Provide an example.

An educated guess about how the natural universe works is called a hypothesis. ‗If I mix one part of hydrogen with one part of oxygen, I can make a substance that contains both the elements.‘ This is made before an experiment.

Easy 68. What are experiments? Why are they used?

Experiments are tests of the natural universe to see if a hypothesis is correct. An experiment to test our previous hypothesis would be to actually mix hydrogen and oxygen and see what happens. Most experiments include observations of small, welldefined parts of the natural universe designed to see results of the experiments.

Moderate 69. What is a theory? Provide an example.

A theory is a general statement that explains a large number of observations. ―All matter is composed of atoms‖ is a general statement, a theory, which explains many observations in chemistry.

Easy 70. Science can be either qualitative or quantitative. Explain with examples.

Qualitative implies a description of the quality of an object. For example, physical properties are generally qualitative descriptions: sulfur is yellow, your math book is heavy, or that statue is pretty. A quantitative description represents the specific amount of something; it means knowing how much of something is present, usually by counting or measuring it.

Moderate Fill in the Blanks 71. _____ is the study of the interactions of matter with other matter and with energy.

Chemistry; Easy 72. Anything that has mass and takes up space can be called _____. matter; Easy 73. A(n) _____ refers to a characteristic that describes matter as it exists.

Physical property; Easy

74. A(n) _____ refers to a change that occurs when a sample of matter changes one or more of its physical properties.

Physical change; Easy 75. A(n) _____ change is the process of demonstrating a chemical property.

Chemical; Easy 76. A(n) _____ refers to a substance that cannot be broken down into simpler chemical substances by ordinary chemical means. element; Easy 77. A(n) _____ is a combination of more than one element and its physical and chemical properties are different from the physical and chemical properties of its constituent elements. compound; easy 78. The constituents of a mixture that are clearly visible to the naked eye, is called a(n) _____. heterogeneous mixture; Easy 79. A combination of two or more substances that is so intimately mixed that the mixture behaves as a single substance is called a(n) _____. homogeneous mixture; Moderate 80. _____ is the process of knowing about the natural universe through observation and experiment.

Science; Easy 81. A(n) _____ refers to an educated guess about how the natural universe works.

Hypothesis; Easy 82. A(n) _____ refers to a test of the natural universe to see if a guess (hypothesis) is correct.

Experiment; Easy 83. A(n) _____ is a general statement that explains a large number of observations. theory; Moderate 84. A(n) _____ refers to a specific statement that is thought to be never violated by the entire natural universe.

Law; Easy 85. A(n) _____ description represents the specific amount of something. quantitative; Easy Chapter 2 Measurements True/False Questions 1. Standard notation is an expression of a number using powers of 10.

False; Easy 2. 0.000411 is an example of using standard notation to express a number.

True; Easy 3. Meter is the SI unit of measuring length of an object.

True; Easy 4. The multiplicative amount of the prefix nano is one thousand.

False; Easy 5. Volume of a vessel is measured using a derived unit.

True; Easy 6. One cubic centimeter is equal to a milliliter.

True; Easy 7. Significant figures represent the limits of what values of a measurement or a calculation we are sure of.

True; Easy 8. Any nonzero digit in a measurement is considered significant.

True; Easy 9. Zeros at the end of a number without a decimal point are considered significant.

False; Easy 10. The number ‗0.006608‘ has six significant figures.

False; Easy 11. The number ‗8.6090 x 103‘ has only four significant figures.

False; Easy 12. The following calculation is performed with the proper number of significant figures. 124 × 10.45 = 1290 True; Moderate 13. As per the 2010 census, the United States population is estimated to be 308.70 million.

This measurement has four significant figures.

False; Moderate 14. If the operations being performed are multiplication or division, the answer has to be limited to the number of significant figures that the data value with the most number of significant figures has.

False; Easy 15. The number 9.666 × 106 has four significant figures.

True; Easy 16. If the expression X+Y=Z is valid, 10X+10Y=10Z is also valid.

True; Easy

17. Conversion factor refers to a fraction that can be used to convert a quantity from one unit to another.

True; Easy 18. 555 nm is equivalent to 0.555 mm.

False; Moderate 19. An exact number is a measured value of an estimate or a calculation.

False; Easy 20. Temperature is a measure of the average amount of static energy a system contains.

False; Easy 21. On the Fahrenheit scale, the freezing point of liquid water is 32°F, and the boiling point of water is 212°F. True; Easy 22. 0°F is equivalent to -32°C. False; Moderate 23. 30°C is the Celsius equivalent of 86°F. True; Moderate 24. Density is a physical property that is defined as a substance‘s mass divided by its volume.

True; Easy 25. A ball made of iron has a mass of 3.78 g. If the density of the ball is 7.87 g/cm3, the volume of the cork is 2.08cm3.

False; Moderate Multiple Choice Questions 26. The straightforward expression of a number is referred to as _____ notation. a. standard b. scientific c. exponential d. technical e. methodical a; Easy 27. _____ notation is an expression of a number using powers of ten. a. Standard b. Customary c. Technical d. Scientific e. General d; Easy 28. Identify the scientific notation of 698. a. 6.98 x 101 b. 6.98 × 102

c. 6.98 x 103 d. 6.98 × 10-2 e. 6.98 × 10-3 b; Easy 29. The exponent‘s value in scientific notation is equal to: a. the power of ten the base number is multiplied by. b. the number of odd decimal digits in the number expressed in standard notation. c. the number of digits after the decimal point in the number expressed in standard notation. d. the number of even decimal digits in the number expressed in standard notation. e. the number of digits in the number expressed in standard notation. a; Easy 30. What does a negative exponent of a number imply? a. The number is imaginary. b. The number‘s value is less than one. c. The number is irrational. d. The number cannot be expressed as a fraction. e. The given number is a negative number. b; Easy 31. The part of a number in scientific notation that is multiplied by a power of 10 is called the _____. a. exponent b. numerator c. association d. coefficient e. denominator d; Easy 32. Which of the following unit systems is commonly used in chemistry? a. FTS system of units b. United States customary units c. International System of Units d. British system of Units e. System of imperial units c; Easy 33. Which of the following is the SI unit of length? a. Meter b. Feet c. Yard d. Inch e. Mile a; Easy 34. Which of the following is a fundamental SI unit? a. Yard b. Kilogram c. Feet

d. Minute e. Joule b; Easy 35. Which of the following is an example of a numerical prefixed SI unit? a. Pound b. Gram c. Yard d. Second e. Centimeter e; Easy 36. Fifteen kilometers is equal to _____. a. 1500 meters b. 150 meters c. 150000 meters d. 15000 meters e. 0.00015 meters d; Easy 37. Which of the following parameters is measured using a derived SI unit? a. Length b. Weight c. Volume d. Time e. Perimeter c; Easy 38. Joule (J) is the unit of energy or work in the International System of Units. A joule can be represented as J = kg×m2÷ s2. Joule is a(n) _____. a. Prefixed unit b. Fundamental unit c. Derived unit d. Imperial unit e. FTS unit c; Easy 39. One liter is _____ of a meter cubed. a. 1/1000th b. 10th c. 100th d. 1/100th e. 1/10th e; Easy 40. Which of the following is measured using a derived SI unit? a. Area of a circular plate b. Perimeter of metal ring c. An individual‘s mass d. Time taken to boil water e. Diameter of a ball

a; Easy 41. The concept of reporting the proper number of digits in a measurement or a calculation is called _____. a. derived units b. prefixed units c. imprecise representation d. loose depiction e. significant figures e; Easy 42. Significant figures represent the limits of _____. a. human ability to approximate the value of a calculation b. what values of a measurement or a calculation we are sure of c. the inaccuracy displayed by a measuring instrument d. the population of data used to arrive at a certain conclusion e. methods used to calculate the accuracy of a calculation b; Easy 43. Which of the following statements is true about the significance of numbers? a. Zeros between non-zero digits are not considered significant. b. All digits of a number are significant, though their relative importance vary. c. Zeros at the end of a number without a decimal point are significant. d. Leading zeroes serve only to put the significant digits in the correct positions. e. Zeros at the end of any number with a decimal point are significant. e; Easy 44. Which of the following zeros is not significant? a. Zeros at the end of a number without a decimal point b. Leading zeros of a decimal number c. Embedded zeros in a decimal number d. Trailing zeros of a decimal number e. Zeros at the beginning of a decimal number b; Easy 45. How many significant figures does the measurement ‗0.00030033‘ have? a. 1 b. 3 c. 4 d. 5 e. 8 d; Moderate 46. Which of the following measurements has 3 significant figures? a. 3670.05 b. 50505 c. 0.06067 d. 0.000060 e. 8.00 x 102 e; Easy

47. For the calculation given below, express the final answer to the proper number of significant figures.

1.23 + 18.788 =? a. 20.018 b. 20.02 c. 20.03 d. 20.029 e. 20.02180 b; Moderate 48. For the calculation given below, express the final answer to the proper number of significant figures.

56.8003 − 50.1 =? a. 6.701 b. 6.702 c. 6.7 d. 6.71 e. 6.7003 c; Moderate 49. For the calculation given below, how many significant figures should the result have?

505 ÷ 5.8030 =? a. 1 b. 2 c. 3 d. 4 e. 5 c; Easy 50. For the below given calculation, express the final answer to the proper number of significant figures.

5.8003 × 20.0=? a. 116 b. 116.1 c. 116.01 d. 116.006 e. 116.0060 a; Moderate 51. How many significant figures does the number ‗3.303×106‘ have? a. 5 b. 6 c. 10 d. 7 e. 4 e; Easy 52. How many feet are there in 15 yards? a. 45 b. 75 c. 5

d. 10 e. 30 a; Easy 53. Expressions used to formally change the unit of a quantity into another unit are called _____. a. natural expressions b. denominators c. dimensions d. conversion factors e. fundamental ratios d; Easy 54. Which of the following terms refers to formally changing the unit of a quantity into another unit? a. Fundamental analysis b. Unit prefixing c. Numerical prefixing d. Signifying numbers e. Factor label method e; Easy 55. Identify the value associated with the prefix milli. a. 1/100 b. 1/1000 c. 100 d. 1000 e. 0.01 b; Easy 56. What is the liter equivalent of 55.12 μL? a. 55120 b. 5512 c. 5.512 × 10−5 d. 5.512× 10−2 e. 5.512× 102 c; Moderate 57. How many cubic centimeters are in 0.1 m3? a. 1000 b. 100 c. 10 d. 10000 e. 100000 e; Moderate 58. Which of the following results is equivalent to 59.9 m/min? a. 599 cm/s b. 5.99 cm/s c. 998 cm/s d. 99.8 cm/s

e. 5990 cm/s d; Moderate 59. Fifteen mm/s equals _____ m/min. a. 9 b. 0.15 c. 0.9 d. 1.5 e. 0.015 c; Moderate 60. 35. How many nanoseconds are in 577.99 ms? a. 5.7799 × 10-8 b. 5.7799 × 102 c. 5.7799 × 105 d. 5.7799 × 108 e. 5.7799 × 109 d; Moderate

61. A(n) _____ number is a number from a defined relationship that technically has an infinite number of significant figures. a. natural b. exact c. complex d. irrational e. fundamental b; Easy 62. A square plot has sides of length 567 cm each. What is the area of the garden plot in square meters? Express your answer in the proper number of significant figures. a. 32.1 b. 22.68 c. 32.148 d. 22.69 e. 32.149 a; Easy 63. _____ is a measure of the average amount of energy of motion, or kinetic energy, a system contains. a. Pressure b. Conductivity c. Temperature d. Density e. Viscosity c; Easy 64. The _____ scale is a temperature scale where 0 degree is the freezing point of water and 100 degree is the boiling point of water. a. Celsius b. Fahrenheit

c. Kelvin d. Newton e. Delisle a; Easy 65. What is 105.4°F in degrees Celsius? a. 38.65°C b. 40.79°C c. 36.24°C d. 39.66°C e. 40.78°C e; Easy 66. The _____ temperature scale uses degrees that are the same size as the Celsius degree, but the numerical scale is shifted up by 273.15 units. a. Rankine b. Fahrenheit c. Kelvin d. Newton e. Delisle c; Easy 67. What is the equivalent of 38°C in Kelvin scale? a. 100.15 K b. 311.15 K c. 300.15 K d. 105.15 K e. 101.15 K b; Moderate 68. Which of the following refers to absolute zero temperature in the Celsius scale? a. 150°C b. 273.15°C c. 295°C d. -273.15°C e. -295°C d; Moderate 69. _____ is a physical property that is defined as a substance‘s mass divided by its volume. a. Dynamic viscosity b. Density c. Kinematic viscosity d. Viscoelasticity e. Pressure b; Easy 70. What is the mass of 40.0 mL of mercury (density 13.6 g/mL)? a. 1436 b. 2449 c. 544 d. 556

e. 2941 c; Easy Essay Questions 71. What is the difference between standard notation and scientific notation?

Standard notation is the straightforward expression of a number. Numbers such as 17, 101.5, and 0.00446 are expressed in standard notation. For relatively small numbers, standard notation is fine. However, for very large numbers, such as 306,000,000, or for very small numbers, such as 0.000000419, standard notation can be cumbersome because of the number of zeros needed to place nonzero numbers in the proper position.

Scientific notation is an expression of a number using powers of 10. Powers of 10 are used to express numbers that have many zeros. It is suitable for very large and very small numbers.

Easy 72. The time taken for a chemical reaction is estimated to be 0.00004 ms. How do you express this measurement, in seconds, using the scientific notation? What are the benefits of doing this?

Here the measurement is expressed using the standard notation. Scientific notation allows you to write very large and small numbers quickly and compactly. This low value measurement can be expressed in scientific notation as 4 × 10-8 s.

Moderate 73. Identify the numbers that are not written in proper scientific notation. Rewrite them, if incorrect, so that they are in proper scientific notation. a. 88.92 × 103 b. 9,943 × 10−5 c. 5.88399 × 105 The first two numbers are not written in proper scientific notation. They can be rewritten as shown below. a. 8.892 × 104 b. 9.943 × 10−2 c. The third number is in proper scientific notation.

Moderate 74. Briefly describe the International System of Units.

The International System of Units or SI system of units specifies certain units for various types of quantities, based on seven fundamental units for various quantities. Some of the important fundamental units are meter for length, seconds for time, and kg (or kilogram) for mass. SI also defines a series of numerical prefixes that refer to multiples or fractions of a fundamental unit to make a unit more conveniently sized for a specific quantity.

Easy 75. What are derived units? Explain by providing an example.

SI system of units allows for derived units based on a fundamental unit or units. There are many derived units used in science. For example, the derived unit for area comes from the idea that area is defined as width times height. Because both width and height are lengths, they both have the fundamental unit of meter, so the unit of area is meter × meter, or meter2 (m2). This is sometimes spoken as ―square meters.‖ A unit with a prefix

can also be used to derive a unit for area, so we can also have cm2, mm2, or km2 as acceptable units for area.

Easy 76. Provide examples of any three multiplicative prefixes and their multiplicative amounts.

Students may list any three of the following.

Prefix

Abbreviation

Multiplicative Amount

gigamegakilodecicentimillimicro nanopico-

G M k d c m μ n p

1,000,000,000 × 1,000,000 × 1,000 × 1/10 × 1/100 × 1/1,000 × 1/1,000,000 × 1/1,000,000,000 × 1/1,000,000,000,000 ×

Easy 77. Explain the concept of significant figures.

The concept of reporting the proper number of digits in a measurement or a calculation is called significant figures. Significant figures (sometimes called significant digits) represent the limits of what values of a measurement or a calculation we are sure of. The convention for a measurement is that the quantity reported should be all known values and the first estimated value.

Easy 78. List the conventions that dictate which numbers in a reported measurement are significant and which are not significant. (a) Any nonzero digit is significant. (b) Any zeros between nonzero digits (i.e., embedded zeros) are significant. (c) Zeros at the end of a number without a decimal point (i.e., trailing zeros) are not significant; they serve only to put the significant digits in the correct positions. However, zeros at the end of any number with a decimal point are significant. (c) Zeros at the beginning of a decimal number (i.e., leading zeros) are not significant; again, they serve only to put the significant digits in the correct positions.

Easy 79. Explain how significant figures are handled in calculations.

Handling of significant figures depends on what type of calculation is being performed. If the calculation is an addition or a subtraction, the rule is as follows: limit the reported answer to the rightmost column that all numbers have significant figures in common. If the operations being performed are multiplication or division, the rule is as follows: limit the answer to the number of significant figures that the data value with the least number of significant figures has.

Easy 80. How are conversions from one unit to another unit of the same type performed?

Conversion factors are fractions that can be used to convert a quantity from one unit to another. These factors are used to coverts units from one unit to another of the same type.

For example, 1 yard = 3 feet. So the conversion factor for feet to yards is 1/3. Feet can be converted to yards by multiplying with 1/3. Yards can be converted to feet by multiplying with 3.

Easy 81. How many millimeters are in 13.66 m? Explain the calculation and creation of conversion factor.

To answer this, one needs to construct a conversion factor between millimeters and meters and apply it correctly to the original quantity. The definition of a millimeter is as follows.

1 mm = 1/1,000 m 1,000 mm =1m The conversion factor = 1000 13.66m = 13.66 ×1000 = 13660mm Moderate 82. What is an exact number? Explain with an example.

An exact number is a number from a defined relationship that technically has an infinite number of significant figures. An exact number is a number from a defined relationship, not a measured one. For example, the prefix kilo means 1,000 exactly, no more or no less.

Easy 83. What is temperature? What is the commonly used temperature scale in the United States?

Temperature is a measure of the average amount of energy of motion, or kinetic energy, a system contains. Temperatures are expressed using scales using units called degrees, and there are several temperature scales in use. In the United States, the commonly used temperature scale is the Fahrenheit scale. On this scale, the freezing point of liquid water (the temperature at which liquid water turns to solid ice) is 32°F, and the boiling point of water (the temperature at which liquid water turns to steam) is 212°F. Easy 84. Explain the Celsius scale. Compare and contrast the Celsius scale with the Kelvin scale.

The Celsius scale (°C) is a temperature scale where 0°C is the freezing point of water and 100°C is the boiling point of water; the scale is divided into 100 divisions between these two landmarks and extended higher and lower.

The fundamental unit of temperature in SI is the kelvin (K). The Kelvin temperature scale uses degrees that are the same size as the Celsius degree, but the numerical scale is shifted up by 273.15 units. That is, the conversion between the Kelvin and Celsius scales is as follows: K = °C + 273.15 °C = K − 273.15 Moderate 85. What is density? How is it measured?

Density is a physical property that is defined as a substance‘s mass divided by its volume:

density =

mass m or d = volume V

Density is usually a measured property of a substance, so its numerical value affects the significant figures in a calculation. Density overall has derived units. Common units for density include g/mL, g/cm3, g/L, kg/L, or even kg/m3.

Moderate Fill in the Blanks 86. 0.00443 is the _____ notation of the number 4.43 × 10-3. standard; Easy 87. A(n) _____ is the raised number to the right of a 10 indicating the number of factors of 10 in the original number. exponent; Easy 88. The part of a number in scientific notation that is multiplied by a power of 10 is called the _____. coefficient; Easy 89. SI specifies certain units for various types of quantities, based on seven _____ units for various quantities. fundamental; Easy 90. 10-9 is the multiplicative amount of the prefix _____. nano; Easy 91. A(n) _____ is a unit that is a product or a quotient of a fundamental unit. derived unit; Easy 92. _____ refers to the limit of the number of places a measurement can be properly expressed with.

Significant figures; Easy 93. Compute and express each answer with the proper number of significant figures.

1.234 + 65.3010 = _____.

66.535; Moderate 94. The number 90,000, when written in scientific notation with significant figures becomes _____.

9 ×104; Moderate 95. A(n) _____ is a fraction that can be used to convert a quantity from one unit to another.

Conversion factor; Easy 96. 5×105 nm is equal to _____ meters.

5×10-4; Easy 97. A(n) _____ is a number from a defined relationship that technically has an infinite number of significant figures.

Exact number; Easy 98. _____ is a measure of the average amount of kinetic energy a system contains.

Temperature; Easy 99. The minimum possible temperature, labeled 0 K, is called _____. absolute zero, Easy 100. _____ is a physical property that is defined as a substance‘s mass divided by its volume.

Density; Easy Chapter 3 Atoms, Molecules, and Ions True/False Questions 1. The smallest piece of an element that maintains the identity of that element is called a molecule.

False; Easy 2. The proton is a heavier subatomic particle than the electron and has a positive charge.

True; Easy 3. The number of protons and electrons in an element‘s nucleus is called the atomic number of the element.

False; Moderate 4. Atoms of the same element with different numbers of electrons are called isotopes.

False; Easy 5. ‗Au‘ is the atomic symbol of the element silver.

False; Easy 6. Uranium-235 isotope of uranium has 143 neutrons.

True; Moderate 7. Hydrogen and oxygen exist naturally as three-atom molecules.

False; Easy 8. The connection between two atoms in a molecule is called a chemical bond.

True; Easy 9. According to chemical nomenclature, SO2 is referred to as sulfur dioxide.

True; Easy 10. The molecule Se2Br2 is called diselenide dibromide.

False; Easy 11. The atomic mass unit is defined as one-twelfth of the mass of a nitrogen-14 atom.

False; Easy 12. The atomic mass of an element is a weighted average of the masses of the isobars of an atom.

False; Easy

13. The molecular mass of dinitrogen trioxide is determined by adding the atomic mass of nitrogen two times with the atomic mass of oxygen three times.

True; Moderate 14. Species with overall negative charges are termed anions.

True; Easy 15. Monatomic ions are formed when an atom loses or gains one or more electrons to attain a charge.

True; Easy 16. I− is an example of a monatomic anion.

True; Easy 17. N3− is called trinitride ion.

False; Easy 18. A proper covalent formula has a cation and an anion in it.

False; Easy 19. In the ionic compound NaCl, sodium ion is the cation and chloride ion is the anion.

True; Moderate 20. According to the nomenclature of ionic compounds, MgCl2 is called magnesium dichloride.

False; Easy 21. Fe2S3 is called iron(II) sulfide.

False; Easy 22. Atoms that contain more than one unit charge are called polyatomic ions.

False; Easy 23. An acid is an ionic compound of the H+ cation dissolved in water.

True; Easy 24. If a compound is composed of hydrogen ions and a polyatomic anion, then the name of the acid is derived from the stem of the polyatomic ion‘s name.

True; Easy 25. HCl is the formula for chloric acid.

False; Easy Multiple Choice Questions 26. The smallest piece of an element that maintains the identity of that element is called a(n) _____. a. compound b. molecule c. atom d. ingredient

e. particle c; Easy 27. Which of the following statements is consistent with the modern atomic theory? a. Atoms of different elements can have the same configuration. b. Atoms combine in whole number ratios to form compounds. c. All matter is composed of isotopes. d. Atoms of the same element are not the same always. e. An electron is smallest piece of an element that maintains its identity. b; Easy 28. Which of the following is a tiny subatomic particle with a negative charge? a. Electron b. Proton c. Neutron d. Photon e. Isotope a; Easy 29. The _____ is a relatively massive subatomic particle with a positive charge. a. isotope b. proton c. neutron d. nucleus e. electron b; Easy 30. Identify the subatomic particle with about the same mass as a proton but no charge. a. Neutron b. Nucleus c. Electron d. Isotope e. Isobar a; Easy 31. What is the approximate mass of a neutron, in kg? a. 6.6 × 10−27 b. 9.1× 10−28 c. 1.6 × 10−27 d. 9.1 × 10−31 e. 1.6 × 10−31 c; Easy 32. The nucleus is made up of what two subatomic particles? a. Proton and electron b. Proton and neutron c. Neutron and electron d. Photon and electron e. Photon and neutron b; Easy

33. Which of the following is an atom that is not a diatomic molecule? a. Helium b. Hydrogen c. Chlorine d. Fluorine e. Oxygen a; Easy 34. The atomic number of an atom refers to the _____. a. number of protons in the atom b. the atomic mass of the nucleus c. the total mass of the electrons d. number of neutrons in the atom e. the atomic mass of the atom a; Easy 35. The sum of the number of protons and neutrons in a nucleus is referred to as the _____ number. a. potential b. elemental c. atomic d. magnetic e. mass e; Easy 36. Two atoms that have the same number of protons with different numbers of neutrons are called _____. a. nuclear isomers b. isotopes c. isotones d. isobars e. chemical isobars b; Easy 37. The number of protons in the nucleus of an atom is 80, while the number of neutrons in the nucleus is 90. What is the mass number of this isotope? a. 10 b. 170 c. 80 d. 90 e. 135 b; Easy 38. Hg is the atomic symbol of the element _____. a. nickel b. helium c. gallium d. mercury e. germanium d; Easy

39. Which of the following is the atomic symbol of gold? a. Gl b. Gd c. An d. Ga e. Au e; Easy 40. Which of the following combinations refer to an isotope? a.

40 20

Ca and 40 K 19

b.

56 26

58 Fe and 26 Fe

c. d.

235 92

U and 235 92 U 8

Li and 4 Be

6 3 235 92

U

e. and 235 91 Pa b; Easy 41. How many neutrons can be found in the atom 11843Tc ? a. 43 b. 161 c. 118 d. 75 e. 32 d; Easy 42. Substances that exist as two or more atoms connected together so strongly that they behave as a single particle are called _____. a. compounds b. polyatoms c. diatomic compounds d. subatomic compounds e. molecules e; Easy 43. Which of the following is not true in regards to sulfur hexafluoride (SF6)? a. It has similar physical properties in the vapor phase as sarin b. It is very easy to detect c. It is not a normal part of the atmosphere d. Used as an aerial tracer for ventilation systems e. It is toxic and highly reactive e; Moderate 44. Hydrogen exists as molecules with _____ atom(s). a. seven b. two c. four d. four

e. five b; Easy 45. A chemical bond is a connection between _____. a. two molecules of a compound b. electrons and the nucleus of an atom c. two atoms in a molecule d. neutrons and protons in the nucleus e. two or more subatomic particles c; Easy 46. What is the molecular formula of elemental hydrogen? a. H2 b. H4 c. H5 d. H6 e. H8 a; Easy 47. Which of the following is name given to SO2 according to chemical nomenclature? a. Sulfur oxide b. Oxo-sulfur c. Sulfur dioxide d. Sulfo-oxide e. Dioxy sulfide c; Easy 48. Which of the following represents diphosphorus hexafluoride? a. (PF3)6 b. P2F4 c. P2FO6 d. P2F6 e. P2O2F6 d; Easy 49. Which of the following constitute a proper molecule? a. Cl b. H2O c. N d. P+ e. Ob; Easy 50. The mass of an electron is _____ atomic mass units. a. 1.00728 b. 0.000549 c. 0.5490 d. 1.879 e. 1.866 b; Easy

51. The _____ of an element is a weighted average of the masses of the isotopes that compose an element. a. atomic number b. mass number c. isotopic mass d. atomic mass e. nuclear mass d; Easy 52. Uranium exists as about 98% U-238 and about 2% U-235. What is the atomic mass of uranium? a. 238.06 b. 237.94 c. 237.14 d. 238.86 e. 239.55 b; Easy 53. The molecular mass is the sum of _____ in a molecule. a. redundant neutrons b. masses of subatomic particles c. masses of protons d. the masses of the atoms e. electronic masses of atoms d; Easy 54. What is the atomic mass of dinitrogen trioxide (atomic mass of oxygen=15.999 u; atomic mass of nitrogen = 14.007 u)? a. 76.011 u b. 30. 010u c. 74.019u d. 160.012u e. 60.019u a; Easy 55. Electrons move from one atom to another to form species with overall electric charges.

Such species are called _____. a. beams b. antiparticles c. ions d. quarks e. composite particles c; Easy 56. Species with overall positive charges are termed _____. a. quarks b. beams c. cations d. anions e. antiparticles c; Easy

57. Which of the following is an example of an anion formed by the gain of two electrons? a. Br− b. Cu2+ c. Na+ d. Se2− e. Rb+ Id; Easy 58. The Cl− ion is formed when a chlorine atom _____. a. loses one electron b. gains one electron c. loses one proton d. gains one proton e. gains one proton and loses one electron b; Easy 59. Which of the following represents the sodium ion? a. Na2+ b. Na+ c. NaOd. Na+OHe. Na2b; Easy 60. Which of the following terms represents the species O2-? a. oxide anion b. oxide cation c. oxygen anion d. oxygen cation e. oxo-cation a; Easy 61. Write the proper ionic formula for the ionic compound obtained from the ions Al3+ and F−. a. Al3F3 b. (AlF)3 c. AlF d. Al3F e. AlF3 e; Easy 62. Ionic compounds form between: a. metals and metals. b. nonmetals and nonmetals c. semiconductors and carbon alloys d. metals and nonmetals. e. semiconductors and nonmetals d; Easy 63. Polyatomic ions are a group of ions that _____.

a. have same mass number and different atomic numbers b. contain more than one atom c. contain more than one positive or negative charge d. have same atomic number and different mass numbers e. contain more than one electron b; Easy 64. Which of the following species is called sodium nitrate? a. NaNO2 b. Na5N c. Na3NOH d. Na2NO e. NaNO3 e; Easy 65. Which of the following polyatomic ions is a cation? a. Bicarbonate b. Carbonate c. Chlorate d. Peroxide e. Ammonium e; Easy 66. Which of the following polyatomic ions is an oxyanion? a. Triiodide b. Cyanide c. Ammonium d. Carbonate e. Calcium ion d; Easy 67. Common table salt (sodium chloride), used on foods, is the ionic compound _____. a. NaCl2 b. SoCl2 c. SoCl d. NaCl e. Na2Cl2 d; Easy 68. An acid is an ionic compound of the _____ dissolved in water. a. K+ cation b. H+ cation c. O- anion d. Na+ cation e. OH- anion b; Easy 69. Which of the following represents acetic acid? a. HC2H3O2 b. H2C2O4 c. HClO4

d. H3PO4 e. H2SO3 a; Easy 70. Which of the following statements is true for all acids? a. Acids do not react with metals. b. Acids do not contain nitrogen. c. Acids have a sour taste. d. Acids do not contain oxygen. e. Acids do not contain water. c; Easy Fill in the Blanks 71. The _____ is a subatomic particle with about the same mass as a proton but no charge. neutron; Easy 72. Atoms of the same element with different numbers of neutrons are called _____. isotopes; Easy 73. The magnitude of the charge is listed as a right _____ next to the symbol of the element. superscript; Easy 74. The smallest part of a substance that has the physical and chemical properties of that substance is called a(n) _____. molecule; Easy 75. The connection between two atoms in a molecule is known as a(n) _____. chemical bond; Easy 76. _____ is a specific system for naming compounds, in which unique substances get unique names.

Chemical nomenclature; Easy 77. The _____ of an element is a weighted average of the masses of the isotopes that compose an element. atomic mass; Easy 78. The _____ is the sum of the masses of the atoms in a molecule. molecular mass; Easy 79. _____ is the proper name for the ionic compound formed between the ions K+ and S2−.

Potassium sulfide; Easy 80. A species with an overall positive charge is called a(n) _____. cation; Easy 81. Species with an overall negative charge are called _____. anions; Easy 82. A compound formed from positive and negative ions is called a(n) _____.

ionic compound; Easy 83. An ion that contains more than one atom is called a(n) _____. polyatomic ion; Easy 84. A(n) _____ is an ionic compound of the H+ cation dissolved in water.

Acid; Easy 85. Hydrochloric acid can be represented by the formula _____.

HCl; Easy Essay Questions 86. Explain the major postulates of modern atomic theory.

Modern atomic theory consists of three parts: 1. All matter is composed of atoms.

2. Atoms of the same element are the same; atoms of different elements are different.

3. Atoms combine in whole number ratios to form compounds.

Easy 87. What is an atom? What are subatomic particles?

The smallest piece of an element that maintains the identity of that element is called an atom. Individual atoms are extremely small. Atoms are composed of smaller parts called subatomic particles. The first part to be discovered was the electron, a tiny subatomic particle with a negative charge. It is often represented as e−. An electron has negative charge. Later, two larger particles were discovered. The proton is a more massive subatomic particle with a positive charge, represented as p+. The neutron is a subatomic particle with about the same mass as a proton but no charge. It is represented as either n or n0.

Easy 88. What are isotopes? Provide two examples.

Atoms of the same element (i.e., atoms with the same number of protons) with different numbers of neutrons are called isotopes. Most naturally occurring elements exist as isotopes. For example, most hydrogen atoms have a single proton in their nucleus.

However, a small number (about one in a million) of hydrogen atoms have a proton and a neutron in their nuclei. Uranium has isotopes such as U-238, U-235, and U-234.

Moderate 89. What is a molecule? Provide examples of elements that exist naturally as molecules.

Many substances exist as two or more atoms connected together so strongly that they behave as a single particle. These multi-atom combinations are called molecules. A molecule is the smallest part of a substance that has the physical and chemical properties of that substance. In some respects, a molecule is similar to an atom. A molecule, however, is composed of more than one atom.

Some elements exist naturally as molecules. For example, hydrogen and oxygen exist as two-atom molecules. Sulfur normally exists as an eight-atom molecule, S8, while phosphorus exists as a four-atom molecule, P4.

Moderate 90. State and explain the rules of chemical nomenclature with an example.

Chemical nomenclature is a very specific system of naming compounds. By following the rules of nomenclature, each and every compound has its own unique name, and each name refers to one and only one compound. The rules of nomenclature for binary molecules are given below.

1. Identify the elements in the molecule from its formula.

2. Begin the name with the element name of the first element. If there is more than one atom of this element in the molecular formula, use a numerical prefix to indicate the number of atoms. Do not use the prefix mono- if there is only one atom of the first element.

3. Name the second element by using three pieces: (a) a numerical prefix indicating the number of atoms of the second element, plus (b) the stem of the element name, plus (3) the suffix -ide.

4. Combine the two words, leaving a space between them.

A molecule whose molecular formula is SO2 is called sulfur dioxide according to these rules.

Moderate

91. What is atomic mass unit? What are the masses of subatomic particles when expressed in this unit?

The atomic mass unit (u) is defined as one-twelfth of the mass of a carbon-12 atom, an isotope of carbon that has six protons and six neutrons in its nucleus. By this scale, the mass of a proton is 1.00728 u, the mass of a neutron is 1.00866 u, and the mass of an electron is 0.000549 u.

Easy 92. Explain how atomic mass of an element is calculated with an example.

The atomic mass of an element is a weighted average of the masses of the isotopes that compose an element. Consider an element that consists of two isotopes, 50% with mass 10 u and 50% with mass 11 u. A weighted average is found by multiplying each mass by its fractional occurrence (in decimal form) and then adding all the products. The sum is the weighted average and serves as the formal atomic mass of the element.

0.50 × 10 u = 5.0 u 0.50 × 11 u = 5.5 u Sum = 10.5 u = the atomic mass of the element.

Moderate 93. What is the molecular mass of a molecule? How is it calculated? Provide an example.

The molecular mass is the sum of the masses of the atoms in a molecule. Although each atom in a molecule is a particular isotope, we use the weighted average, or atomic mass, for each atom in the molecule.

For example, in order to determine the molecular mass of dinitrogen trioxide, N2O3, we would need to add the atomic mass of nitrogen two times with the atomic mass of oxygen three times: 2 N masses = 2 × 14.007 u = 28.014 u 3 O masses = 3 × 15.999 u = 47.997 u Total = 76.011 u = the molecular mass of N2O3 Moderate 94. What are ions? What are the two types of ions?

Electrons can move from one atom to another; when they do, species with overall electric charges are formed. Such species are called ions. Species with overall positive charges are termed cations, while species with overall negative charges are called anions.

Easy 95. What are ionic compounds? Provide two examples.

Compounds formed from positive and negative ions are called ionic compounds.

Individual atoms can gain or lose electrons. Consider the ionic compound between Na+ and Cl−. Each ion has a single charge, one positive and one negative. The ionic compound formed here is NaCl. The ionic compound between magnesium cations (Mg2+) and oxide anions (O2−) is MgO. Moderate 96. How are monatomic ions formed? Provide examples.

Individual atoms can gain or lose electrons. When they do, they become monatomic ions.

When atoms gain or lose electrons, they usually gain or lose a characteristic number of electrons, and so take on a characteristic overall charge. H+, Na+, and K+ are examples of cations. O2−, S2− and Cl- are examples of anions.

Easy 97. Explain the formation of an ionic compound between Mg2+ ions and Cl− ions.

For the ionic compound between Mg2+ ions and Cl− ions, we now consider the fact that the charges have different magnitudes, 2+ on the magnesium ion and 1− on the chloride ion. To balance the charges with the lowest number of ions possible, we need to have two chloride ions to balance the charge on the one magnesium ion. The compound formed is represented as MgCl2.

Moderate 98. How do you name binary molecular compounds? Explain with two examples.

Begin the name with the element name of the first element. If there is more than one atom of this element in the molecular formula, use a numerical prefix to indicate the number of atoms. Do not use the prefix mono- if there is only one atom of the first element. Name the second element by using three pieces: a numerical prefix indicating the number of atoms of the second element, plus the stem of the element name (e.g., ox for oxygen, chlor for chlorine, etc.), plus the suffix -ide. Combine the two words, leaving a space between them. Examples: SF6 (named sulfur hexachloride), and P2O5 (diphosphorus pentaoxide) Moderate 99. Explain the terms polyatomic ions and oxyanions.

A group of ions that contain more than one atom are called polyatomic ions. Most of them also contain oxygen atoms, so sometimes they are referred to as oxyanions.

CH3COO−, the acetate ion, is a polyatomic ion and an oxyanion.

Easy 100. What are acids? Explain their nomenclature.

Acid is an ionic compound of the H+ cation dissolved in water. Acids have their own nomenclature system. If an acid is composed of only hydrogen and one other element, the name is hydro- + the stem of other element + -ic acid. If a compound is composed of hydrogen ions and a polyatomic anion, then the name of the acid is derived from the stem of the polyatomic ion‘s name. Typically, if the anion name ends in -ate, the name of the

acid is the stem of the anion name plus -ic acid; if the related anion‘s name ends in -ite, the name of the corresponding acid is the stem of the anion name plus -ous acid.

Moderate

Chapter 4 Chemical Reactions and Equations True/False Questions 1. The equation, H2 + O2  H2O, is not balanced.

True; Easy 2. CO2 is a reactant in the chemical equation 2C2H6 + 7O2  4CO2 + 6H2O. False; Easy 3. The chemical equation CH4 + O2  CO2 + H2O cannot be balanced to form a proper equation.

False; Easy 4. A decomposition reaction generates more than one product from a single reactant.

True; Easy 5. Iodide is an example of a halogen.

False; Easy 6. When replacing cations in single displacement reactions, chemical reactivity trends can be predicted simply by using their relative positions on the periodic table.

False; Easy 7. Iron can replace aluminum by undergoing a single-replacement reaction.

False; Moderate 8. Calcium can react with LiNO3 to replace Li and form calcium nitrate.

False; Easy 9. A double-replacement reaction occurs when parts of two ionic compounds are exchanged to make two new compounds.

True; Easy 10. A precipitation reaction occurs when two ionic compounds are dissolved in water and form a new ionic compound that does not dissolve.

True; Easy 11. Li3PO4 cannot be dissolved in water.

False; Moderate 12. Lead (II) carbonate (PbCO3) cannot be dissolved in water.

True; Moderate 13. Generally, all compounds of NO3− and C2H3O2− dissolve in water.

True; Easy 14. CaCl2 dissociates to form one Ca2+ ions and one Cl2- ion.

False; Easy 15. A decomposition reaction starts from a single substance and produces more than one substance.

True; Easy 16. The reaction, C3H8  C3H4 + 2H2, is an example of a combustion reaction False; Easy 17. H3O+ is called trihydroxo ion.

False; Easy 18. Acids and bases neutralize to produce a salt compound and water.

True; Easy 19. Any ionic compound that is formed from a reaction between an acid and a base is called a salt.

True; Easy 20. Neutralization reactions do not proceed if one of the reactants is not in the aqueous phase.

False; Easy 21. Reduction is defined as the gain of one or more electrons by an atom.

True; Easy 22. Oxidation is defined as the gain of one or more protons by an atom.

False; Easy 23. Oxidation and reduction always occur together in reality.

True; Easy 24. Atoms in their elemental state are assigned an oxidation number of 1.

False; Easy 25. The oxidation number of oxygen molecule (O2) is -2.

False; Easy Multiple Choice Questions 26. Consider the equation, 2H2 + O2 → 2H2O. Here, H2 and O2 are called _____. a. reactants b. products c. catalysts d. modules e. alkalis a; Easy

27. In chemical equations, the number of atoms of each element in the reactants must be the same as the number of atoms of each element in the products. This principle is formed in accordance with the _____. a. law of identity b. second law of thermodynamics c. law of conservation of matter d. law of attraction of materials e. third law of thermodynamics c; Easy 28. An unbalanced chemical equation can be balanced by _____. a. changing the number of atoms in the product b. adding oxygen molecules in either sides of the equation c. changing the mass number of the molecules that react or are produced d. changing the number of molecules that react or are produced e. adding hydrogen ions in either sides of the equation d; Easy 29. Which of the following is the balanced chemical equation that represents nitrogen and hydrogen reacting to produce ammonia? a. N2 + 2H3  2NH3 b. N2 + 3H2  2NH3 c. N3 + 7H2 + O2  3NH3 + 2H2O + H+ d. N3 + 5H2  3NH3 + H+ e. N2 + 3H2  NH3 +H3N b; Easy 30. Which of the following is an example of a single-replacement reaction? a. HCl (aq) + NaOH (aq)  NaCl (aq) + H2O b. AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq) c. 2HCl(aq) + Zn(s)  ZnCl2(aq) + H2(g) d. Ba(OH)2 (s) + 2 CuCNS (s)  Ba(CNS)2 (aq) + CuOH (s) e. HCH3COO (aq) + NaHCO3 (s)  NaCH3COO (aq) + CO2 (g) + H2O (l) c; Easy 31. 2HCl (aq) + Zn(s)  ? a. ZnCl + HCl b. ZnCl2 + H2 c. ZnCl + H2Cl d. H2Zn + Cl2 e. HZnCl2 + H+ b; Easy 32. Which of the following elements is a halogen? a. Iodine b. Helium c. Neon d. Hydrogen e. Lithium a; Easy

33. Identify the products if a double-replacement reaction occurs here. 2AgNO3 + ZnCl2 →? a. Zn(NO)2 + 2Ag + 2Cl2 b. Zn(NO)2 + 2AgCl2 c. Ag2Cl + Zn(NO3)2 d. 2AgCl + Zn(NO3)2 e. Ag2Zn + N2 + 3O2 d; Moderate 34. Which of the following elements is likely to replace potassium (K) through a singlereplacement reaction? a. Lithium b. Calcium c. Sodium d. Magnesium e. Aluminum a; Easy 35. Identify an element that cannot replace iron through a single-replacement reaction. a. Potassium b. Lithium c. Calcium d. Copper e. Aluminum d; Easy 36. Identify the chemical reaction for which no reaction can be predicted. a. FeCl2 + Zn → ? b. Fe + Cu(NO3)2 → ? c. FeCl2 + Cu → ? d. Cl2 + 2NaBr → ? e. AlPO4 + Mg → ? c; Easy 37. Identify a double-replacement reaction from the following. a. Cu + 2AgNO3 → 2Ag + Cu(NO3)2 b. Fe + Cu(NO3)2 → Fe(NO3)2 + Cu c. Ca + 2H2O → Ca(OH)2 + H2 d. Zn + 2HCl → ZnCl2 + H2 e. CuCl2 + 2AgNO3 → Cu(NO3)2 + 2AgCl e; Easy 38. A _____ reaction occurs when two ionic compounds are dissolved in water and form a new ionic compound that does not dissolve. a. precipitation b. contamination c. condensation d. diffusion e. dispersion a; Easy

39. _____ rules are general statements that predict which ionic compounds dissolve and which do not. a. Composition b. Dissociation c. Solubility d. Combustibility e. Decomposition c; Easy 40. Which of the following compounds dissolves in water? a. AgCl b. CaCO3 c. HgCl2 d. PbI2 e. Ba(OH)2 e; Easy 41. Which of the following compounds is not likely to dissolve in water? a. HCl b. AgCl c. NaCl d. KI e. RbBr b; Easy 42. Which of the following carbonates can be dissolved in water? a. CaCO3 b. FeCO3 c. NiCO3 d. (NH4)2CO3 e. Al2(CO3)3 d; Easy 43. Which of the following is the precipitate in the possible double-replacement reaction between Na2SO4 and SrCl2? a. Strontium sulfate b. Sodium hydride c. Sodium chloride d. Strontium carbonate e. Sodium oxide a; Moderate 44. When ionic compounds dissolve, the ions physically separate from each other. This process is called _____. a. composition b. dissociation c. diffusion d. combustion e. decomposition b; Easy

45. Which of the following is not an ionic compound? a. CaCO3 b. NH3 c. NaCl d. KBr e. CuS b; Easy 46. How many ions are formed when MgSO4 dissociates? a. 2 b. 3 c. 4 d. 5 e. 7 a; Easy 47. Which of the following ions are formed when AgNO3 dissociate? a. N2b. N3c. AgN2d. O2e. NO3e; Easy 48. Consider the reaction NaCl + AgNO3. Which of the following is a spectator ion in this ionic reaction? a. Cl− b. Ag+ c. Na+ d. N3e. O2c; Easy 49. A chemical equation with the spectator ions removed is called a(n) _____ equation. a. spectator ionic b. balanced c. unbalanced d. complete ionic e. net ionic e; Easy 50. An ion that does nothing in the overall course of a chemical reaction is called a(n) _____ ion. a. balanced b. neutral c. nonaligned d. spectator e. intermediate d; Easy 51. A _____ reaction produces a single substance from multiple reactants.

a. synthesis b. decomposition c. neutralization d. combustion e. diffusion a; Easy 52. Which of the following is an example of a synthesis reaction? a. 2NaHCO3 → Na2CO3 + CO2 + H2O b. BaSO4 → Ba2+ + SO42− c. 2H2 + O2 → 2H2O d. C3H8 + 5O2 → 3CO2 + 4H2O e. 4NH3 + 3O2 → 2N2 + 6H2O c; Easy 53. A _____ reaction starts from a single substance and produces more than one substance. a. synthesis b. diffusion c. combustion d. neutralization e. decomposition e; Easy 54. Which of the following chemical equations represents a decomposition reaction? a. C2H5OH + 3O2 → 2CO2 + 3H2O b. 2HBr + Cl2 → 2HCl + Br2 c. Na2CO3 →Na2O + CO2 d. H2 + Cl2 → 2HCl e. Na2O + CO2 → Na2CO3 c; Easy 55. A(n) _____ reaction occurs when a reactant combines with oxygen, many times from the atmosphere, to produce oxides of all other elements as products. a. synthesis b. decomposition c. neutralization d. combustion e. diffusion d; Easy 56. Which of the following is an example of a combustion reaction? a. 2HBr + Cl2 → 2HCl + Br2 b. C2H5OH + 3O2 → 2CO2 + 3H2O c. 2CO2 + 3H2O → C2H5OH + 3O2 d. Na2CO3 →Na2O + CO2 e. Na2O + CO2 → Na2CO3 b; Easy 57. Which of the following is a common product of combustion reactions? a. H2 b. CO2

c. O2 d. N3 e. CH4 b; Easy 58. Identify a product obtained from the combustion reaction of ammonia. a. N2 b. O2 c. CO2 d. H2O e. H2 d; Easy 59. A(n) _____ is any compound that increases the amount of hydrogen ion (H+) in an aqueous solution. a. metal b. salt c. acid d. base e. mineral c; Easy 60. A base is a compound that increases the amount of _____ ion in an aqueous solution. a. oxide b. hydroxide c. hydrogen d. hydride e. hydronium b; Easy 61. Which of the following represents a hydronium ion? a. H3O+ b. H+ c. H2+ d. He. OHa; Easy 62. The reaction of an acid and a base is called a _____ reaction. a. synthesis b. combustion c. decomposition d. synthesis e. neutralization e; Easy 63. Which of the following terms refer to an ionic compound that is formed from a reaction between an acid and a base? a. salt b. polymer c. alkali

d. soap e. hydroxide a; Easy 64. The chemical reaction represented by the equation HCl + KOH  H2O + KCl is an example of a(n) _____ reaction. a. combustion b. neutralization c. synthesis d. decomposition e. composition b; Easy 65. The process of losing one or more electrons by an atom is called _____. a. neutralization b. dehydration c. oxidation d. reduction e. synthesis c; Easy 66. Chemical reactions that involve the transfer of electrons are called _____ reactions. a. synthesis b. combustion c. composition d. neutralization e. redox e; Easy 67. Which of the following rules is valid when assigning oxidation numbers to atoms? a. Atoms in their elemental state are assigned an oxidation number of 1. b. Atoms in the form of molecules have an oxidation number of 1. c. Hydrogen is assigned a +1 oxidation number when it exists as the hydride ion d. Atoms in monatomic ions are assigned an oxidation number equal to their charge. e. Oxygen and hydrogen are usually assigned an oxidation number of zero. d; Easy 68. In the compound calcium chloride, what is the oxidation number given to chloride? a. -1 b. 1 c. -2 d. 2 e. 0 a; Easy 69. When an oxidation number of an atom is decreased in the course of a reaction, that atom is being _____. a. synthesized b. neutralized c. oxidized d. decomposed

e. reduced e; Easy 70. Consider the chemical reaction, Fe + CuSO4 → FeSO4 + Cu. Which of the following is being oxidized here? a. SO4 b. O c. S d. Cu e. Fe e; Easy Fill in the Blanks 71. Initial substances in a chemical equation are called _____. reactants; Easy 72. Final substances in a chemical equation are called _____. products; Easy 73. _____ is a chemical reaction in which one element is substituted for another element in a compound.

Single-replacement reaction; Easy 74. _____ are general statements that predict which ionic compounds dissolve and which do not.

Solubility rules; Easy 75. _____ is the process of an ionic compound separating into ions when it dissolves.

Dissociation; Easy 76. A chemical equation in which the dissolved ionic compounds are written as separated ions is called a(n) _____. complete ionic equation; Easy 77. The _____ lists the elements that will replace elements below them in single replacement reactions. activity series; Easy 78. A chemical reaction in which a single substance is produced from multiple reactants is called a(n) _____. composition reaction; Easy 79. A(n) _____ is a chemical reaction in which a reactant combines with oxygen to produce oxides of all other elements as products.

Combustion reaction; Easy 80. A compound that increases the amount of H+ ions in an aqueous solution is called a(n) _____. acid; Easy

81. A(n) _____ is a compound that increases the amount of OH− ions in an aqueous solution.

Base; Easy 82. The _____ are the elements in the next-to-last column on the periodic table.

Halogens; Easy 83. Any ionic compound that is formed from a reaction between an acid and a base is called a(n) _____. salt; Easy 84. The loss of one or more electrons by an atom is called _____. oxidation; Easy 85. _____ is a number assigned to an atom that helps keep track of the number of electrons on the atom.

Oxidation number; Easy Essay Questions 86. What is a chemical equation? Explain with an example.

A chemical equation is a concise way of representing a chemical reaction. Consider the following equation, H2 + O2  2H2O. The initial substances are called reactants, and the final substances are called products. Chemical equations should be balanced. When the reactants and products of a chemical equation have the same number of atoms of all elements present, we say that the equation is balanced.

Easy 87. Balance the following equation and explain the steps. CH4 + O2 → CO2 + H2O When the reactants and products of a chemical equation have the same number of atoms of all elements present, we say that an equation is balanced. We have two oxygen atoms on the left and three on the right. This equation can be balanced by multiplying the oxygen molecule on the left side by two and the water molecule on the right by two.

After balancing the equation becomes, CH4 + 2O2 → CO2 + 2H2O Moderate 88. What is a single-replacement reaction? Explain with an example.

A single-replacement reaction is a chemical reaction in which one element is substituted for another element in a compound, generating a new element and a new compound as products. Consider the following equation, 2HCl + Zn  ZnCl2 + H2 The hydrogen atoms in HCl are replaced by Zn atoms, and in the process a new element, hydrogen, is formed.

Easy 89. Write a short note on predicting the chemical reactivity trends in single-replacement reactions.

Chemical reactivity trends can be predicted when replacing anions in simple ionic compounds by simply using their relative positions on the periodic table. However, when replacing the cations, the trends are not as straightforward. A list called the activity series can be used in these situations. It lists the elements that will replace elements below them in single-replacement reactions.

Easy 90. Finish the following equation, balance, then state the type of reaction.

Cu2SO4 + Fe(NO3)2  A double-replacement reaction occurs when parts of two ionic compounds are exchanged, making two new compounds. A characteristic of a double-replacement equation is that there are two compounds as reactants and two different compounds as products. There are two equivalent ways of considering a double-replacement equation: either the cations are swapped, or the anions are swapped. Cu2SO4 + Fe(NO3)2  2 CuNO3 + FeSO4 Moderate 91. What are solubility rules? Provide two examples of general solubility rules.

Solubility rules are general statements that predict which ionic compounds dissolve and which do not. Some examples of solubility rules are given below.

(1) All compounds of Li+, Na+, K+, Rb+, Cs+, and NH4+ dissolve in water.

(2) All compounds of NO3− and C2H3O2− dissolve in water.

(3) All compounds of CO32− and PO43− except those have Li+, Na+, K+, Rb+, Cs+, or NH4+ do not dissolve in water.

Easy 92. How do chemical reactions that include ionic compounds differ from chemical reactions that include molecular compounds?

One important aspect about ionic compounds that differs from molecular compounds has to do with dissolving in a liquid, such as water. When molecular compounds, such as sugar, dissolve in water, the individual molecules drift apart from each other. When ionic compounds dissolve, the ions physically separate from each other.

Moderate 93. Explain the following: (1) complete ionic equation, (2) spectator ions, (3) net ionic equation.

A complete ionic equation is a chemical equation in which the dissolved ionic compounds are written as separated ions. Some of the ions in a complete ionic equation do nothing in the overall course of the chemical reaction. They are present, but they do not participate in the overall chemistry. Such ions are called spectator ions. What remains when the spectator ions are removed is called the net ionic equation, which represents the actual chemical change occurring between the ionic compounds.

Easy 94. What is a composition reaction? Provide an example.

A composition reaction (sometimes also called a combination reaction or a synthesis reaction) produces a single substance from multiple reactants. A single substance as a product is the key characteristic of the composition reaction. There may be a coefficient other than one for the substance, but if the reaction has only a single substance as a product, it can be called a composition reaction. The following reaction is an example of a composition reaction.

2H2O + O2  2H2O Easy 95. Explain how a combustion reaction occurs. Provide an example.

A combustion reaction occurs when a reactant combines with oxygen, many times from the atmosphere, to produce oxides of all other elements as products; any nitrogen in the reactant is converted to elemental nitrogen, N2. Many reactants, called fuels, contain mostly carbon and hydrogen atoms, reacting with oxygen to produce CO2 and H2O. For example, the balanced chemical equation for the combustion of methane, CH4, is as follows: CH4 + 2O2  CO2 + 2H2O Easy 96. What are acids and bases?

An acid is any compound that increases the amount of hydrogen ion (H+) in an aqueous solution. The chemical opposite of an acid is a base. The equivalent definition of a base is that a base is a compound that increases the amount of hydroxide ion (OH−) in an aqueous solution.

Easy 97. Explain neutralization reaction.

The reaction of an acid and a base is called a neutralization reaction. Although acids and bases have their own unique chemistries, the acid and base cancel each other‘s chemistry to produce a rather innocuous substance—water. The general reaction between an acid and a base is: acid + base  water + salt.

The term salt is generally used to define any ionic compound (soluble or insoluble) that is formed from a reaction between an acid and a base. Neutralization reactions are one type of chemical reaction that proceeds even if one reactant is not in the aqueous phase.

Easy 98. Compare and contrast oxidation and reduction.

In reality, electrons are lost by some atoms and gained by other atoms simultaneously.

Oxidation is defined as the loss of one or more electrons by an atom. Reduction is defined as the gain of one or more electrons by an atom. In reality, oxidation and reduction always occur together. Chemical reactions that involve the transfer of electrons are called oxidation-reduction reactions.

Moderate 99. List the rules for assigning oxidation numbers to atoms.

The rules for assigning oxidation numbers to atoms are as follows: (1) Atoms in their elemental state are assigned an oxidation number of 0.

(2) Atoms in monatomic (i.e., one-atom) ions are assigned an oxidation number equal to their charge. Oxidation numbers are usually written with the sign first, then the magnitude, which differentiates them from charges.

(3) In compounds, fluorine is assigned a −1 oxidation number; oxygen is usually assigned a −2 oxidation number (except in peroxide compounds [where it is −1] and in binary compounds with fluorine [where it is positive]); and hydrogen is usually assigned a +1 oxidation number (except when it exists as the hydride ion, H−, in which case rule 2 prevails).

(4) In compounds, all other atoms are assigned an oxidation number so that the sum of the oxidation numbers on all the atoms in the species equals the charge on the species (which is zero if the species is neutral).

Easy 100. When do you say an atom is oxidized or reduced?

When an oxidation number of an atom is increased in the course of a redox reaction, that atom is being oxidized. When an oxidation number of an atom is decreased in the course of a redox reaction, that atom is being reduced. Oxidation and reduction are thus also defined in terms of increasing or decreasing oxidation numbers, respectively.

Easy Chapter 5 Stoichiometry and the Mole True/False Questions 1. 4Fe + 3O2 → 2Fe2O3. This equation indicates that eighteen molecules of oxygen is needed to produce twelve molecules of iron(III) oxide.

True; Easy 2. One gram of carbon-12 contains one mole of carbon atoms.

False; Easy 3. The numerical value of things in a mole is often called Avogadro‘s number.

True; Easy 4. A mole of a substance has the same mass in grams as one unit (atom or molecules) has in atomic mass units.

True; Easy 5. The molar mass of carbon is 12 g. This indicates that 12 × 6.022 × 1023 atoms are present in one gram of carbon.

False; Easy 6. Consider the equation, 2Na + Cl2 → 2NaCl. This equation shows that 6.022 × 1023 units of NaCl can be produced when one mole of Cl2 and two moles of Na react.

False; Easy 7. If the molar mass of HCl is 36.46, the mass of 15 mol of HCl is 546.9 g.

True; Easy 8. If the density of water is 1 g/mL, 18 moles are present in 18 ml of water.

False; Easy 9. 100H2 + 50O2  100H2O this equation is balanced though it is not conventional.

True; Easy 10. Fe2O3 + Fe + 6 HCl → 3 FeCl2 + 3 H2O. Here, 120 moles of Fe2O3 can be used to produce 360 moles of FeCl2 True; Easy 11. 2C4H10 + 13O2  8CO2 + 10H2O. This equation indicates that two moles of C4H10 and thirteen moles of O2 is needed for the reaction to be balanced.

True; Easy 12. Fe2O3 + Fe + 6 HCl → 3 FeCl2 + 3 H2O. This chemical equation is balanced in terms of moles.

True; Easy 13. Consider the chemical reaction, 2Al + 3Cl2  2AlCl3. Here, 56 g of Al can produce 276.54 g of AlCl3 (Molar masses: Al = 26.98 g; Cl = 35.45 g).

False; Moderate 14. Consider the chemical reaction, 2Al + 3Cl2  2AlCl3. Here, 212.7 g of chlorine can produce a maximum of 266.66 g of AlCl3 (Molar masses: Al = 26.98 g; Cl = 35.45 g).

True; Moderate 15. Mass quantities of one substance can be related to mass quantities using a balanced chemical equation.

True; Easy 16. When calculating the amount of product assuming that all the reactant reacts, we are calculating the actual yield.

False; Easy 17. The percent yield can be calculated using the formula, percent yield =actual yield÷ theoretical yield × 100%.

True; Easy 18. If percent yield determined is greater than 90%, it indicates an error.

False; Easy 19. The theoretical yield of a reaction is 50 g and the actual yield is 40 g. The percent yield of this reaction is 80%.

True; Moderate 20. Consider the reaction, N2 + 3H2  2NH3. 15 g of NH3 is formed from 18 g of N2. The percent yield in this example is 83.33% (Molar masses: N = 14 g; H = 1.0 g).

False; Moderate 21. The percent yield of the various stages in a reaction is given below.

A→ Intermediary1 (percent yield = 70%) → Intermediary2 (percent yield = 80%) → B (Percent Yield = 90%). The overall percent yield of this reaction is 50.%.

True; Easy 22. The reactant that is available in abundance during a chemical reaction is called excess reagent.

True; Easy 23. 4As + 3O2  2AsO3. 50 g As and 28 g O2 react with each other. Oxygen (O2) is the limiting reagent in this chemical reaction (Molar masses: As = 79.92 g; O = 16 g).

False; Moderate 24. The reactant gives the least amount of a particular product in a chemical reaction can be called the limiting reagent for that product.

True; Easy

25. CH4 + 2 O2 → CO2 + 2 H2O. 50 g of CH4 and 120 g of O2 react with each other. The reactant O2 is in excess here. (Molar masses: C = 12 g; H = 1 g; O = 16 g) False; Moderate Multiple Choice Questions 26. The relating of one chemical substance to another using a balanced chemical reaction is called _____. a. computational chemistry b. informatics c. analytical chemistry d. stoichiometry e. stereochemistry d; Easy 27. How many molecules of oxygen should react with hydrogen molecules to form 14 molecules of water? a. 3 b. 4 c. 6 d. 7 e. 8 d; Easy 28. Na + Cl2 → NaCl. This is an unbalanced chemical equation showing the formation of NaCl. How many chlorine molecules are needed to form 252 molecules of NaCl? a. 28 b. 42 c. 63 d. 126 e. 252 d; Easy 29. C8H18 + O2 → CO2 + H2O. Balance this equation and identify the number H2O molecules formed when 6 molecules of C8H18 react with 75 molecules of oxygen. a. 54 b. 48 c. 36 d. 32 e. 18 a; Moderate 30. Fe2O3 (s) + 3SO3 (g)  Fe2(SO4)3. This equation explains the formation of iron(III) sulfate. How many molecules of Fe2(SO4)3 are formed if 144 molecules of SO3 gas is passed through a beaker containing iron(III) oxide? a. 288 b. 144 c. 72 d. 24 e. 48 e; Easy

31. How many moles of atoms can be found in 70 grams of carbon-12? a. 0.5 b. 1 c. 60 d. 6 e. 5 d; Easy 32. One _____ refers to the number of things equal to the number of atoms in exactly 12 g of carbon-12. a. quantum b. mole c. Faraday d. Henry e. carbon unit b; Easy 33. Avogadro‘s number refers to _____. a. the numerical value of things in a mole b. the total number of atoms in a kg of a compound c. a ratio of the masses of reactants and the masses of all the products d. the total number of neutrons in a molecule e. the total number of protons in a compound a; Easy 34. One mole equals _____ things. a. 1.6726 × 1027 b. 6.022 × 1023 c. 3.16 × 1024 d. 2.303 × 1024 e. 1.903 × 1024 b; Easy 35. How many molecules are present in 3.16 mol of CaCl2? a. 3.16 × 1023 b. 6.16 × 1023 c. 3.16 × 1024 d. 6.03 × 1023 e. 1.90 × 1024 e; Easy 36. How many kilograms of carbon-12 would contain 6.000 × 1026 atoms? a. 12.35 b. 1.65 c. 11.95 d. 10.65 e. 0.12 c; Easy 37. What is the mass of 1 mole of diatomic hydrogen?

a. 2.6 g b. 2.0158 g c. 12 g d. 2.9 g e. 1.0079 g b; Easy 38. The molar mass of NaCl is 58.45 g. How many molecules does a gram of NaCl contain? a. 1.030 × 1022 b. 2.015 × 1023 c. 6.022 × 1023 d. 1.03 × 1023 e. 6.022 × 1022 a; Easy 39. Consider the chemical reaction. 2H2 +O2 → 2H2O. How many water molecules can be formed from 15.0 g H2 if O2 is available abundantly (Molar mass of H2 = 2.016 g)? a. 4.48 × 1023 b. 5.36 × 1023 c. 1.67 × 1023 d. 4.48 × 1024 e. 5.36 × 1024 d; Moderate 40. How many CO molecules are formed when one mole of CO2 reacts with 1.000 mole of carbon? a. 6.022 × 1023 b. 6.022 × 1024 c. 1.204 × 1024 d. 3.022 × 1023 e. 2.409 × 1024 c; Moderate 41. What is the mass of 3.56 mole of CuCl2 if the molar mass of CuCl2 is 134.45 g/mol? a. 479 b. 478.642 c. 478.6 d. 478 e. 480 a; Moderate 42. Consider the chemical equation P4 + 5O2  P4O10. How many moles of oxygen are needed to convert 4 moles of phosphorous to P4O10? a. 4 b. 5 c. 10 d. 15 e. 20 e; Easy

43. Consider the chemical reaction 2CuCl2 + SO2 + 2H2O → 2CuCl + 2HCl + H2SO4. How many grams of water is required to produce 14.0 moles of sulphuric acid? a. 14.0 b. 28.0 c. 144 d. 504 e. 505 d; Moderate 44. How many moles of oxygen would yield 120 moles of water molecules in the following chemical reaction?

2C4H10 + 13O2  8CO2 + 10H2O a. 78 b. 104 c. 160 d. 156 e. 130 c; Easy 45. Consider the unbalanced equation H2 + O2 → H2O. How many molecules of water is involved in the equation if it is interpreted in terms of moles? a. 1.2044 × 1024 b. 6.022 × 1023 c. 1.2044 × 1023 d. 6.022 × 1024 e. 3.2 × 1027 a; Easy 46. Sixteen moles of aluminium (Al) react with twenty-four moles of chlorine. How many moles of AlCl3 are obtained as a result of this reaction? a. 12 b. 16 c. 20 d. 24 e. 28 b; Easy 47. Consider the chemical equation, Pd + 2 CuCl2 → 2 CuCl + PdCl2. How many moles of CuCl2 is needed to produce 1.6 moles of PdCl2? a. 0.8 b. 1 c. 1.2 d. 1.6 e. 3.2 e; Easy 48. Consider the unbalanced chemical reaction Al + Cl2  AlCl3. How many moles of AlCl3 can be obtained from 360 grams of Cl2? (Molar masses: Cl = 35.45 g/mol, Al = 26.98 g/mol) a. 70.9 b. 400

c. 270 d. 130 e. 450 (these answers do not follow significant figures) e; Hard 49. How many molecules are present in 58 grams of CuCl2 if the molar mass of copper is 63.55 g and molar mass of chlorine is 35.45 g? a. 1.67 b. 2.64 c. 0.32 d. 0.43 e. 0.59 d; Easy 50. Which of the following is the conversion factor for AlCl3 and HCl considering the following equation?

2AlCl3 + 3H2O  Al2O3 + 6HCl a. 1/1 b. 1/3 c. 1/2 d. 1/4 e. 1/6 b; Easy 51. How many moles of HCl will be produced when 30.000g of H2O are reacted according to the chemical equation 2AlCl3 + 3H2O  Al2O3 + 6HCl? (Molar masses: H2O = 18.015 g; HCl = 36.46 g) a. 72.92 b. 36.03 c. 121.43 d. 109.38 e. 145.84 c; Easy 52. Consider the chemical equation, N2 + 3H2  2NH3. How many grams of nitrogen are needed to produce 50.00 grams of ammonia? (Molar masses: N2 = 28.013 g; NH3 = 17.03 g) a. 56.026 b. 41.12 c. 34.06 d. 51.09 e. 82.25 b; Easy 53. How many grams of SO3 can be produced from 15.00 grams of oxygen molecule when reactions takes place according to the following equation 2SO2 + O2  2SO3? (Molar masses: S = 32.07 g; O = 16.00 g) a. 75.07 b. 80.07 c. 64.00

d. 63.00 e. 150.13 a; Easy 54. The amount that is really produced in a reaction is called the _____. a. molar yield b. molar mass c. percent yield d. limiting reagent e. actual yield e; Easy 55. Which of the following is an assumption made when calculating theoretical yield? a. Chemical reaction is conducted in vacuum b. All the reactant in a chemical reaction reacts with others. c. Chemical reaction involves heating process d. Room temperature is maintained throughout the reaction e. Actual yield is more than the theoretical yield. b; Easy 56. Theoretical yield of a reaction is 30 g where as the actual yield is only 15 g. What is the percent yield for this reaction? a. 45% b. 50% c. 22.5% d. 90% e. 60% b; Easy 57. The percent yield for a reaction is estimated to be 70.0%. Calculate the actual yield if the theoretical yield is 105 g of product. a. 52.5 g b. 70 g c. 73.5 g d. 105 g e. 102.5 g c; Easy 58. Which of the following cannot be a valid percent yield value? a. 0% b. 30% c. 60% d. 90% e. 120% e; Easy 59. Consider the chemical equation, 2H2+O2 → 2H2O (Molar masses: H2 = 2.016 g; O2 = 32 g; H2O = 18.016 g). If the percent yield of this reaction is estimated to be only 90%, calculate the amount of water that can be produced from 90.00 g oxygen. a. 91.2 b. 72.06

c. 202.66 d. 182.4 e. 36.03 d; Moderate 60. Consider the chemical reaction, H2CO3 → H2O + CO2. Here, the percent yield is calculated to be 60%. Calculate the amount of H2CO3 needed to produce 56.00 g CO2 (Molar masses: H2CO3 = 62.03 g; H2O = 18.015 g; CO2 = 44.00 g). a. 62.03 g b. 93.045 g c. 72.06 g d. 131.58 g e. 78.95 g d; Hard 61. In a chemical reaction A and B reacts to become E. C and D are the intermediaries of this reaction. The percent yields of the reaction are given below. What is the overall percent yield?

A+B→C (Percent yield=70%) C→D (Percent yield=90%) D→E(Percent yield=50%) a. 210% b. 50% c. 90% d. 70% e. 31.5% e; Easy 62. In a chemical reaction, A and B reacts to form D. C is the intermediate product of this reaction. Percent yield of the reaction A+B→C is 50.% and percent yield of the reaction C→D is 90.%. How many moles of D can be produced when 14 moles of A and B react? a. 5.6 b. 7 c. 19.6 d. 11.2 e. 39.2 a; Easy 63. Which of the following can be used as a measure to compare the actual yield and theoretical yield? a. Molar yield b. Mass number c. Limiting reagent d. Percent yield e. Nuclear mass d; Easy 64. Consider the chemical equation 2H2 + O2 →2H2O. Which of the following is the limiting reagent when 18 g hydrogen reacts with 30 g oxygen? (Molar masses: H = 1.008 g; O = 16.00 g) a. Hydrogen

b. Water c. Oxygen d. H+ ion e. OH- ion c; Easy 65. A limiting reagent is _____. a. the reactant that controls the atmospheric conditions limit b. a catalyst used in chemical reaction c. the reactant which gives the lesser amount of product in a chemical reaction d. a negative catalyst that slows down chemical reactions e. the reactant that behaves according to atmospheric conditions c; Easy 66. Reactant that _____ is called the excess reagent. a. gives the least amount of product b. is available in excess c. has the largest mass number d. has the smallest mass number e. has the smallest atomic number b; Easy 67. Consider the reaction, 4As + 3O2  2AsO3. You have 149.84g of As to perform the reaction. O2 will be considered a limiting reagent if the available quantity is less than _____ g. (Molar masses As = 74.92 g; O = 16.00 g) a. 2.000 b. 74.92 c. 1.000 d. 122.92 e. 149.84 d; Easy 68. Consider a chemical reaction that follows the equation NaHCO3 + CH3COOH → NaCH3COO + CO2 + H2O. 80 g of NaHCO3 and 65 g CH3COOH is available for performing this reaction. Which of the following is considered to be in excess?

(Molar masses: NaHCO3 = 84 g; CH3COOH = 60.05 g) a. NaCH3COO b. NaHCO3 c. CO2 d. H2O e. CH3COOH e; Easy 69. Consider the equation, 2Rb + MgCl2  Mg + 2RbCl. If 82.00 g of Rb are available, what is the amount of MgCl2 needed? (Molar masses: Rb = 85.47 g; Mg = 24.3 g; Cl = 35.45 g) a. 95.2 b. 85.47 c. 45.67 d. 82.47

e. 91.33 c; Easy 70. How many grams of C2H2 can be formed when one gram of carbon reacts with one gram of hydrogen according to the reaction 2C + H2  C2H2? (Molar masses: C = 12 g; H = 1 g) a. 6.32 b. 1.08 c. 2.16 d. 4.08 e. 12.6 b; Easy Essay Questions 71. What is stoichiometry? Why is it used?

The relating of one chemical substance to another using a balanced chemical reaction is called stoichiometry. Using stoichiometry is a fundamental skill in chemistry. It greatly broadens one‘s ability to predict what will occur and, more importantly, how much is produced.

Easy 72. Briefly explain the concept of a mole and molar mass.

A mole is a number of things equal to the number of atoms in exactly 12 g of carbon-12.

A mole equals 6.02214179 × 1023 things. Molar masses of substances refer to the mass for 1 mol of things.

Easy 73. How many moles would represent 1.5 × 1026 molecules of oxygen? Explain your answer.

One mole contains 6.022× 1023 molecules.

Therefore 1.5 × 1026 molecules of oxygen = 1.5 × 1026 ÷ 6.022 × 1023 = 249 moles= 250 moles (significant figures) Moderate 74. Are all balanced equations conventional? Explain your answer.

The chemical equation is balanced as long as the coefficients are in the ratio of a conventional balanced equation. The conventional balanced equation uses the lowest possible coefficients. Consider the equation, 100H2 + 50O2  100H2O. This equation is balanced, but not conventional. The conventional balanced equation is 2H2 + O2  2H2O. Moderate 75. A + 3B → 2C. Interpret this chemical equation in terms of moles.

One mole of A reacts with 3 moles of B to form 2 moles of C. 1mole A + 3 moles B → 2 moles C Easy 76. Compare and contrast mole-mole calculation and mole-mass calculation.

A stoichiometry calculation when one starts with moles of one substance and convert to moles of another substance using the balanced chemical equation. Mole to mass calculation is a calculation in which you start with a given number of moles of a

substance and calculate the mass of another substance involved in the chemical equation, or vice versa.

Moderate 77. Explain mole-mass calculation with an example.

Consider the equation, 2Al + 3Cl2  2AlCl3. Suppose we have 100 g of Cl2. Chemical equations are not balanced in terms of grams; they are balanced in terms of moles. So to use the balanced chemical equation to relate an amount of Cl2 to an amount of AlCl3, we need to convert the given amount of Cl2 into moles. We know how to do this by simply using the molar mass of Cl2 as a conversion factor.

100g chlorine equals 100/70.9 = 1.41 moles.

1.41 mol chlorine * 2 mol AlCl3 ÷ 3 mol Cl2 = 0.94 moles 100 g chlorine can produce 0.94 moles of AlCl3. That is, it can produce 125.34 g AlCl3.

Easy 78. Consider the equation, 2AlCl3 + 3H2O  Al2O3 + 6HCl. How many moles of Al2Cl3 can be produced from 1 kg of AlCl3? Explain the answer.

(Molar masses: AlCl3 = 133.34 g/mol; Al2O3 = 101.96 g/mol) 1 kg AlCl3 equals 7.45 moles of AlCl3.

7.45 moles of AlCl3 × 1 mol Al2O3 ÷ 2 mol AlCl3 = 3.725 moles.

1 kg AlCl3 can yield 3.725 moles of Al2O3.

Moderate 79. What is mass-mass calculation?

It is a calculation in which you start with a given mass of a substance and calculate the mass of another substance involved in the chemical equation.

Easy 80. Explain the concepts of theoretical yield and actual yield.

When we calculate an amount of product assuming that all the reactant reacts, we calculate the theoretical yield, an amount that is theoretically produced as calculated using the balanced chemical reaction. In many cases, less of a product is made during the course of a chemical reaction. The amount that is actually produced in a reaction is called the actual yield.

Easy 81. What is percent yield?

The percent yield is a comparison between the actual yield and the theoretical yield and is defined as

percent yield =

actual yield ´ 100 theoretical yield Percent yield always has units of percent.

Proper percent yields are between 0% and 100%—again, if percent yield is greater than 100%, an error has been made.

Easy 82. N2 + 3H2 2NH3 A synthesis produces 50 g ammonia from 60 g N2. What is the theoretical yield and the percent yield?

1 mole of nitrogen should have produced 2 moles of ammonia.

I.e., 60 g nitrogen should have produced 72.857 g ammonia

Theoretical yield = 72.857 g Actual yield = 50 g Percent yield = 50 ÷ 72.857 × 100 = 68.63%.

Moderate 83. What is limiting reagent? Provide an example.

The reactant you run out of during a chemical reaction is called the limiting reagent.

Consider the reaction 2H2 + O2 →2H2O 4 g of hydrogen can react with 32 g oxygen to produce 36 g H2O. If 16 g oxygen and 4 g hydrogen react, only 18 g H2O can be produced. 2 g of hydrogen remains un-reacted. Thus oxygen is called the limiting reagent in this reaction.

Easy 84. When do you say a reactant is in excess? Explain with an example.

Consider the reaction 2H2 + O2 →2H2O 4 g of hydrogen can react with 32 g oxygen to produce 36 g H2O. If 16 g oxygen and 4 g hydrogen react, only 18 g H2O can be produced. 2 g of hydrogen remains un-reacted. Thus oxygen is called the limiting reagent in this reaction.

The reactant you run out of during a chemical reaction is called the limiting reagent. The other reactant or reactants are considered to be in excess. A crucial skill in evaluating the conditions of a chemical process is to determine which reactant is the limiting reagent and which is in excess.

Easy 85. Consider the equation, 2NaCl + H2SO4 → Na2SO4 + 2HCl. Which reactant is the limiting reagent when 50.00 g NaCl and 60.00 g H2SO4 react? How much is in excess?

50.00 g NaCl equals 50/58.44 = 0.8560 moles 60.00 g H2SO4 equals 60/98.08 = 0.6117 Here, H2SO4 is in excess because 0.856 moles of NaCl can react with only 0.428moles of H2SO4. NaCl is the limiting reagent. To find excess, take the limiting reagent and find the number of grams for the excess reagent. 50.00g of NaCl/58.44= 0.8560 moles 0.8560 moles/2 moles= .4280 moles of H2SO4 .4280 moles H2SO4 * 98.08 g= 41.98 grams of H2SO4 60.00 g H2SO4 – 41.98g = 18.04 grams of H2SO4 in excess Moderate Fill in the Blanks 86. _____ molecules of H2O are formed when 16 molecules of hydrogen react with 8 molecules of oxygen.

Sixteen; Easy 87. A(n) _____ is a number of things equal to the number of atoms in exactly 12 g of carbon12. mole; Easy 88. Avogadro‘s number refers to the _____. numerical value of things in a mole; Easy 89. The molecular mass of CaCO3 is _____ if the molar masses of Ca, C, and O are 40, 12, and 16 g/mol, respectively.

100; Easy 90. Six grams of carbon-12 contain _____ atoms.

3.011 × 1023; Easy 91. In the chemical reaction 2H2 + O2  2H2O, in order to produce 50. moles of H2O, 50. moles of hydrogen should react with _____ moles of oxygen molecules.

25; Easy 92. 2H2 + O2  2H2O. This chemical reaction when represented in terms of moles becomes _____.

2 mol H2 + 1 mol O2 → 2 mol H2O; Moderate 93. A(n) _____ is a stoichiometry calculation when one starts with moles of one substance and convert to moles of another substance using the balanced chemical equation.

Mole-mole calculation; Easy 94. _____ grams of NH3 will be produced when 50.00 g of N2 are reacted according to the following chemical equation. N2 + 3H2  2NH3 (Molar masses: N→14; H→1) 60.71; Easy 95. A(n) _____ is a calculation in which you start with a given number of moles of a substance and calculate the mass of another substance involved in the chemical equation, or vice versa.

Mole-mass calculation; Easy 96. A(n) _____ is a calculation in which you start with a given mass of a substance and calculate the mass of another substance involved in the chemical equation.

Mass-mass calculation; Easy 97. The _____ is a comparison between the actual yield and the theoretical yield. percent yield; Easy 98. The actual yield of a reaction is 67 g whereas the theoretical yield is 78 g. The percent yield for this reaction is _____.

86%; Easy 99. The reactant you run out of during a chemical reaction is called the _____. limiting reagent; Easy 100. One gram of H2 and one gram of O2 react to form water. The limiting reagent in this example is _____.

O2; Easy

Chapter 6 Gases True/False Questions 86. According to the kinetic theory of gases, the size of a gas particle is large compared to the distances that separate them.

False; Easy 87. An ideal gas is a gas that exactly follows the statements of the kinetic theory.

True; Easy 88. The SI unit of pressure is atmosphere (atm).

False; Easy 89. Pascal (Pa) is equal to 1 N/m2.

True; Easy 90. One torr is an equivalent unit of one millimeter of mercury True; Easy 91. P1T1 = P2T2, represents the Charles‘s law of gases.

False; Easy 92. Pressure and volume of a gas are inversely related, if other parameters are kept the same.

True; Easy 93. Robert Boyle invented the mercury barometer.

False; Easy 94. When applying Charles‘s law, amount of gas and volume are considered to be constants.

False; Easy 95. A sample of a gas has an initial volume of 15.0 L and an initial temperature of 235 K. The temperature of the gas increases to 313 K, if its volume increases to 20.0 L. True; Moderate 96. A sample of a gas has an initial volume of 16.8 L and an initial temperature of 50°C. The volume of the gas increases to 33.6 L if its temperature is increased to 100° C. False; Moderate 97. Gay-Lussac‘s law relates volume of a gas with absolute temperature.

False; Easy P1 P2  98. According to Gay-Lussac‘s law, T1 T2 at constant V and n.

True; Easy

99. Avogadro‘s law states that equal volumes of different gases at the same temperature and pressure contain the same amount of gas.

True; Easy 100.

If a 2.45 L volume of gas contains 4.5 × 1021 gas particles, a 4.9 L volume of gas will contain 4.5 × 1022 gas particles.

False; Moderate 101. variable.

False; Easy

The combined gas law considers the amount of gas as a

102.

False; Easy

Real gases are ideal.

103.

True; Easy

The equation, PV = nRT is called the ideal gas law.

104.

A 6.00 mol sample of O2 has a pressure of 0.80 atm and a temperature of 68°C. Its volume is 41.8L. False; Easy 105.

True; Easy

All gases occupy the same volume at STP.

106.

False; Easy

2 moles of H2 are present in 4 L at STP.

107.

True; Easy

The density of O2 is 1.54 g/L at 30°C and 1.20 atm.

108.

If mixture of H2 at 5.00 atm and N2 at 8.00 atm is stored in a container, the mixture‘s pressure will be 10.1 atm.

False; Easy 109. temperature.

True; Easy

The vapor pressure of water increases with an increase in its

110.

The mole fraction of He is 0.43 in a mixture containing He at 0.60 atm and Ne at 0.80 atm.

True; Easy Multiple Choice Questions

111.

Which of the following statements is consistent with the kinetic theory of gases? f. Gases consist of tiny stationary particles of matter. g. Gases constantly collide with each other. h. The size of a gas particle is large when compared to the distances between them. i. A net loss of energy occurs when particles collide. j. Strong interactive forces exist between the particles of a gas. b; Easy 112.

A gas that exactly follows the statements of the kinetic theory is called _____. a. real gas b. kinetic gas c. dynamic gas d. ideal gas e. stationary gas d; Easy 113.

X is a gas that deviates slightly from agreeing perfectly with the kinetic theory of gases. X is an example of _____ gas. a. real b. ideal c. supreme d. static e. sturdy a; Easy 114. _____ refers to the force of all the gas particle/wall collisions divided by the area of the wall. a. Temperature b. Pressure c. Buoyancy d. Density e. Concentration b; Moderate 115.

If a force of 1.50 x103 N is pressed against an area of 3.00 m2, what is the pressure in torr? a. 3.75 b. 5.65 c. 500 d. 565 e. 3,800 a; Moderate

116.

Which of the following is the SI approved unit of pressure? a. Pounds/inch b. Millimeters of mercury c. Atmosphere d. N/cm2 e. Pascal e; Easy 117.

Which of the following is equivalent to one kilopascal? a. 760 torr b. 1.00 atm c. 10 N/m2 d. 1000 Pa e. 14.7 lb/in2. d; Easy 118.

1 atm is equal to _____ torr. a. 540 b. 760 c. 916 d. 512 e. 144 b; Easy 119.

Pressure at a point deep in an ocean is estimated to be 1.074 × 103 kPa. What is the average atmospheric pressure at that point? a. 1.074 × 10-1 atm b. 1.060 × 106 atm c. 1.060 x 101 atm d. 1.074 × 10-4 atm e. 1.074 × 103 atm c; Moderate 120.

How many atmospheres are there in 2,326 torr? a. 2.370 b. 23.26 c. 11.63 d. 3.060 e. 1.160 d; Moderate 121.

Which of the following statements is true about the relationship between pressure and volume? a. Pressure and volume are independent of each other. b. Pressure and volume decrease with an increase in temperature. c. Pressure increases as volume decreases.

d. Pressure and volume are not related to temperature. e. Pressure increases as volume increases. c; Easy 122.

Boyle‘s law explains the relationship between _____. a. pressure and temperature b. volume and temperature c. the amount of gas and temperature d. pressure and the amount of gas e. pressure and volume e; Easy 123.

Experiments show that the volume of a gas is related to its temperature in _____. a. Celsius b. Fahrenheit c. Kelvin d. Joule e. Newton c; Easy 124.

Which of the following properties of gases must be held constant when applying Boyle‘s law? a. Pressure b. Amount of gas c. Volume d. Shape of the container e. Size of the container b; Easy 125.

A gas has an initial pressure of 1.823 atm and an initial volume of 8.026 L. What is its new pressure if volume is changed to 6.323 L? Assume temperature and amount of gas are held constant. a. 0.9342 atm b. 1.436 atm c. 3.105 atm d. 2.314 atm e. 27.84 atm d; Easy 126.

A gas has an initial pressure of 1,080 torr and an initial volume of 586 mL. What is its new volume if pressure is changed to 1.30 atm? Assume temperature and amount of gas are held constant. a. 370 mL b. 486 L c. 641 L

d. 486 L e. 641 mL e; Moderate 127.

Which of the following laws states

V  constant for a given amount of a T

gas, when pressure is constant? a. Charles‘s law b. Boyle‘s law c. Avogadro‘s law d. Graham‘s law e. Gay-Lussac‘s law a; Easy 128.

What is the Kelvin equivalent of 145°C? a. 752 b. 293 c. 418 d. 116 e. 292 c; Easy 129.

A sample of gas has an initial volume of 5.00 L and an initial temperature of 360. K. What is the new volume if the temperature is increased to 900. K? Assume constant pressure and amount for the gas. a. 2.00 L b. 3.00 L c. 8.00 L d. 12.5 L e. 5.00 L d; Easy 130.

A sample of gas has an initial volume of 87.01 mL and an initial temperature of 295.5 K. What is the new temperature if the volume is decreased to 50.00 mL? Assume constant pressure and amount for the gas. a. 169.8 K b. 514.2 K c. 158.6 K d. 530.5 K e. 14.72 K a; Easy 131.

A sample of gas has a volume of 60. mL and an initial temperature of 380 K. The temperature must be changed to _____ in order to halve the volume.

Assume constant pressure and amount for the gas. a. 273 K

b. 76 K c. 190 K d. 760 K e. 780 K c; Easy 132.

A gas has an initial volume of 65.08 mL and an initial temperature of 577°C. What is its new volume if temperature is changed to 417°C? Assume pressure and amount of gas are held constant. a. 56.2 mL b. 47.0 mL c. 80.2 mL d. 90.1 mL e. 52.8 mL e; Easy 133.

Which of the following gas laws relates pressure with absolute temperature? a. Avogadro‘s law b. Graham‘s law c. Charles‘s law d. Boyle‘s law e. Gay-Lussac‘s law e; Easy 134.

What is the final pressure of a gas whose initial pressure is 590 torr, initial temperature is 308 K, and final temperature is 286 K? Assume volume and amount of gas are held constant. a. 585 torr b. 386 torr c. 150 torr d. 550 torr e. 631 torr d; Easy 135.

What is the final temperature of a gas whose initial pressure is 455 torr, initial temperature is 105°C, and final pressure is 690 torr? Assume volume and amount of gas are held constant. a. 300.°C b. 69°C c. 159°C d. –23.7°C e. 186°C a; Easy 136.

Which of the following laws consider amount of gas as a variable?

a. Charles‘s law b. Avogadro‘s law c. Combined gas law d. Boyle‘s law e. Gay-Lussac‘s law b; Easy 137.

A 3.8 mL volume of gas contains 5.6 × 1022 gas particles. How many gas particles are there in 15 mL if the gas is at the same pressure and temperature? a. 3.8 × 1022 b. 1.5 × 1023 c. 2.2 × 1023 d. 1.5 × 1022 e. 5.4 × 1022 c; Easy 138.

Which of the following laws proposes a relationship between pressure, temperature, and volume of a gas when the amount of gas is a constant? a. Gay-Lussac‘s law b. Boyle‘s law c. Combined gas law d. Charles‘s law e. Avogadro‘s law c; Easy 139.

Which of the following is a constant when applying the combined gas law? a. Pressure b. Volume c. Amount of gas d. Temperature e. Density c; Easy 140.

A sample of gas reduces its pressure and absolute temperature by 55 percent. What is the percentage change in volume? a. 0 b. 20 c. 40 d. 50 e. 100 a; Easy 141.

A sample of gas at an initial volume of 65.0 mL, an initial pressure of 866 torr, and an initial temperature of 10.0°C simultaneously changes its temperature to 60.0°C and its volume to 46.0 mL. What is the final pressure of the gas?

a. 154 torr b. 2,150 torr c. 6,210 torr d. 168 torr e. 1,440 torr e; Easy 142.

Which of the following equations represent the ideal gas law? a. P1V1 = P2V2 V1 V2  T T2 1 b.

P1 P2  T T2 1 c.

PV PV 1 1  2 2 T2 d. T1 e. PV = nRT e; Easy 143. a. b. c. d.

Which of the following statements is true about the ideal gas law?

Pressure is considered to be a constant when applying the ideal gas law.

The ideal gas law does not require a change in the conditions of a gas sample.

Volume is considered to be a constant when applying the ideal gas law.

According to ideal gas law, temperature is not an independent property of gases. e. The ideal gas law is represented by the equation P1V1 = P2V2. b; Easy

144.

A 5.86 mol sample of Ne has a pressure of 263 torr and a temperature of 54.0°C. What is its volume? a. 0.598 L b. 455 L c. 0.0988 L d. 460 L e. 58.6 L b; Moderate 145.

A 1.6608 mol sample of O2 has a pressure of 0.9561 atm and a temperature of 386 K. What is its volume? a. 55.0 L b. 59.0 L c. 30.2 L d. 50.6 L e. 59.8 L a; Easy

146.

For a 0.0962 mol sample of O2, V = 58 L and T = 456 K. What is its pressure? a. 0.61 atm b. 0.021 atm c. 0.21 atm d. 0.047 atm e. 0.062 atm e; Easy 147.

Which of the following is considered the standard temperature under STP? a. 0 K b. 273 K c. 160 K d. 300 K e. 313 K b; Easy 148.

What is the molar volume of oxygen gas (O2) at STP? a. 44.8 L b. 22.4 L c. 103.04 L d. 411.2 L e. 716.8 L b; Easy 149.

How many moles of H2 are present in 69.2 L at STP? a. 17.3 mol b. 34.6 mol c. 3.09 mol d. 4.00 mol e. 8.65 mol c; Easy 150.

Which of the following statements is consistent with the Dalton‘s law of partial pressures? a. The total pressure of a mixture of gases is less than the sum of the partial pressures of individual gases. b. The total pressure of a gas mixture is equal to the sum of the partial pressures of the components. c. The total pressure of a gas mixture is calculated by assessing the weighted average of the molecular pressures of the components. d. The total pressure of a gas mixture is independent of the partial pressures of the components. e. The total pressure of a mixture of gases is more than the sum of the partial pressures of individual gases. b; Moderate

151.

A mixture of H2 at 566 torr and O2 at 600 torr is in a container. What is the total pressure in the container? a. 3,396 torr b. 566 torr c. 600 torr d. 1,166 torr e. 3,000 torr d; Easy 152.

A mixture of helium at 1.600 atm and neon at 900.0 torr is in a container.

What is the total pressure in the container? a. 1,008 torr b. 901.6 torr c. 450.8 torr d. 2,116 torr e. 598 torr d; Easy 153.

H2 stored in a 50.0 L container has a pressure of 500.0 torr. O2 stored in a 25.0 L container has a pressure of 800.0 torr. If both these gases are released together to be stored in a 100.0 L container, what is the total pressure in the 100.0 L container? a. 800 torr b. 500 torr c. 450 torr d. 650 torr e. 1,300 torr c; Easy 154.

CO2 generated by the decomposition of ZnCO3, is collected in a 5.6 L container over water. If the temperature is 40.0°C and the total pressure inside the container is 580 torr, how many moles of CO2 were generated? a. 0.09 mol b. 0.74 mol c. 0.18 mol d. 0.15 mol e. 0.86 mol d; Moderate 155.

A container has a mixture of H2 at 1.2 atm and O2 at 1.6 atm. What is the mole fraction of H2? a. 0.55 b. 0.45 c. 0.57 d. 0.40

e. 0.43 e; Easy Essay Questions 156.

What are the postulates of the kinetic theory of gases? a) Gases consist of tiny particles of matter that are in constant motion. b) Gas particles are constantly colliding with each other and the walls of a container.

These collisions are elastic; that is, there is no net loss of energy from the collisions. c) Gas particles are separated by large distances, with the size of a gas particle tiny compared to the distances that separate them. d) There are no interactive forces (i.e., attraction or repulsion) between the particles of a gas. e) The average speed of gas particles is dependent on the temperature of the gas.

Easy 157.

What are ideal gases? Can real world gases be called ideal gases?

An ideal gas is a gas that exactly follows the statements of the kinetic theory.

Real gases are not ideal. Many gases deviate slightly from agreeing perfectly with the kinetic theory of gases. However, most gases adhere to the statements so well that the kinetic theory of gases is well accepted by the scientific community.

Easy 158.

What is pressure? How is it calculated?

Pressure is the force of all the gas particle/wall collisions divided by the area of the wall. Pressure is calculated as a ratio of force and area.

Easy 159.

What are the various units of pressure? Show how to convert one pressure unit into another using dimensional analysis.

The formal, SI-approved unit of pressure is the pascal (Pa), which is defined as 1 N/m2. A common unit of pressure is the atmosphere (atm), which was originally defined as the average atmospheric pressure at sea level. A more reliable and common unit is millimeters of mercury (mmHg), which is the amount of pressure exerted by a column of mercury exactly 1 mm high. An equivalent unit is the torr, which equals 1 mmHg. Example: 1.03 atm 760 torrs= 783 torrs 1.00 atm Easy 160.

Briefly explain Boyle‘s law.

Boyle‘s law is a gas law that relates pressure and volume at constant temperature and amount.

Itstates that that pressure and volume are inversely related.

P × V = constant, at constant n and T.

This can be rewritten as P1V1 = P2V2 at constant n and T. Easy 161.

Briefly explain Charles‘s law.

If the temperature of a gas is expressed in kelvins, then experiments show that the V ratio of volume to temperature is a constant: = constant.

T V V This equation can be rewritten as 1  2 . These conditions remain valid as long T1 T2 as pressure and the amount of the gas remain constant. This gas law is commonly referred to as Charles‘s law.

Easy 162.

Briefly describe Gay-Lussac‘s law. Highlight the difference between GayLussac‘s and Charles‘s law.

Gay-Lussac‘s law relates pressure with absolute temperature. In terms of two sets of data, Gay-Lussac‘s law is, P1 P2  at constant V and n.

T1 T2 This law has a structure very similar to that of Charles‘s law, only with different variables—pressure instead of volume.

Easy 163.

What is Avogadro‘s law? Explain.

Avogadro‘s law states that equal volumes of different gases at the same temperature and pressure contain the same number of particles of gas. Because the number of particles is related to the number of moles (1 mol = 6.022 × 1023 particles), Avogadro‘s law essentially states that equal volumes of different gases at the same temperature and pressure contain the same amount (moles, particles) of gas. Put mathematically into a gas law, Avogadro‘s law is, V1 V2  at constant n n1 n2 Easy 164.

What is the combined gas law?

Combined gas law is a gas law that combines pressure, volume, and temperature.

This gas law is known as the combined gas law, and its mathematical form is, PV PV 1 1  2 2 at constant n.

T1 T2 Easy 165.

Explain the concept of ideal gas law.

The equation PV = nRT is called the ideal gas law. It relates the four independent properties of a gas at any time. The constant R is called the ideal gas law constant.

Its value depends on the units used to express pressure and volume.

Easy 166.

Explain the significance and use of the ideal gas law.

The ideal gas law is used like any other gas law, with attention paid to the unit and making sure that temperature is expressed in Kelvin. However, the ideal gas law does not require a change in the conditions of a gas sample. The ideal gas law implies that if you know any three of the physical properties of a gas, you can calculate the fourth property.

Easy 167.

Explain the concept of standard temperature and pressure.

Standard temperature and pressure (STP) is defined as exactly 100 kPa of pressure (0.986 atm) and 273 K (0°C). For simplicity, 1 atm is used as the standard pressure. Defining STP allows us to compare more directly the properties of gases that differ from each other.

Easy 168.

What is molar volume? What is the molar volume of gases at STP The molar volume is the volume of 1 mol of a gas. It is a property shared among all gases. At STP, the molar volume of a gas can be easily determined by using the ideal gas law: L.atm   (1 atm)V = (1 mol)  0.08205  (273 K) mol.k   All the units cancel except for L, the unit of volume. So V = 22.4 L Easy 169.

Briefly explain Dalton‘s law of partial pressures.

Dalton‘s law of partial pressures states that the total pressure of a gas mixture, Ptot, is equal to the sum of the partial pressures of the components, Pi. Ptot  P1  P2  P3  ...   Pi # of gases

Easy

170.

What is mole fraction?

The mole fraction, χi, is the ratio of the number of moles of component i in a mixture divided by the total number of moles in the sample: moles of component i ci = total number of moles

Easy

Fill in the Blanks 171.

The _____ theory of gases is a fundamental model that describes the physical properties of gases. kinetic; Easy 172.

A gas that exactly follows the statements of the kinetic theory is called a(n) _____. ideal gas; Easy 173. _____ is defined as the force of all the gas particle/wall collisions divided by the area of the wall Pressure; Easy 174.

15 torr is equal to _____ kilopascals. two; Easy 175.

A _____ is a simple mathematical formula that allows you to model or predict the behavior of a gas. gas law; Easy 176. _____ law states that P × V = constant at constant n and T. Boyle’s; Easy 177.

Charles‘s law states that _____ is a constant at constant P and n.

V ; Easy T 178.

The Celsius scale equivalent of 248K is _____ °C. -25; Easy 179. _____ law relates pressure with absolute temperature.

Gay-Lussac’s; Easy 180. _____ law states that equal volumes of different gases at the same temperature and pressure contain the same number of particles of gas.

Avogadro’s; Easy 181.

The gas law that unites pressure, volume, and temperature is called the _____. combined gas law; Easy 182.

Value of the ideal gas constant in terms of joules/molK is _____.

8.314; Easy

183. _____ is defined as exactly 100 kPa of pressure (0.986 atm) and 273 K (0°C).

Standard temperature and pressure (STP); Easy 184.

The _____ of a gas is the pressure that an individual gas in a mixture has. partial pressure; Easy 185. The ratio of the number of moles of one of the components in a mixture divided by the total number of moles in the sample is called that component‘s _____. mole fraction; Easy

Chapter 7 Energy and Chemistry True/False Questions 186.

One newton-meter is a unit used to measure both work and energy, it is also equal to one joule.

True; Easy 187.

False; Easy

One calorie is equal to 4.184 kJ.

188.

An isolated system is a system that does not allow a transfer of energy or matter into or out of the system.

True; Easy 189. heated or cooled.

False; Easy

The total energy of an isolated system increases when it is internally

190.

The amount of work done by a gas when expanding can be determined by multiplying the final volume and external pressure.

False; Easy 191.

The work performed by a gas is 507 J when it expands from 56.3 L to 58.8 L against an external pressure of 2.0 atm.

True; Easy 192.

For a given object, the amount of heat is inversely proportional to the mass of the object.

False; Easy 193.

The heat involved when 150.0 g of Fe increases its temperature from 300.°C to 350.°C is 3370 J. The specific heat of Fe is 0.449 J/g·°C. True; Easy

194.

Enthalpy change refers to the pressure of a container when a heat change takes place.

False; Easy 195.

A chemical equation that includes an enthalpy change is called a thermochemical equation.

True; Easy 196. reaction.

True; Easy

A chemical reaction that has a positive ΔH is an endothermic

197. thermometry.

False; Easy

The process of measuring changes in enthalpy is called

198.

Given the thermochemical equation.

2H2(g) + O2(g) → 2H2O H = -566 kJ This equation shows that two moles of hydrogen is needed to produce two moles of H2O. True; Easy 199.

Given the thermochemical equation.

N2(g) + 3H2(g) → 2NH3(g) ΔH = −91.8 kJ 91.8 kJ of energy is given off when one mole of N2 reacts.

True; Easy 200.

Given the thermochemical equation N2(g) + O2(g) → 2NO(g) ΔH = +180.6 kJ 180.6 kJ of energy is released when two moles of NO are produced.

False; Easy 201.

Given the thermochemical equation.

N2(g) + O2(g) → 2NO(g) ΔH = +180.6 kJ If 7.00g of N2 will be reacted, 45.2 kJ will be added. (Molar masses: N = 14 g; O = 16 g) True; Easy 202.

Given the thermochemical equation.

2C(s) + 2O2(g) → 2CO2(g) ΔH = −393.50 kJ The ΔH for the reaction CO2 → C + O2 is +196.75 kJ. True; Easy 203.

Hess‘s law states that when chemical equations are combined algebraically, their enthalpies can be combined in exactly the same way.

True; Easy 204.

True; Easy

If a chemical reaction is reversed, the sign on ΔH is changed.

205.

If the ΔH for N2 + O2 → 2NO is +180 kJ, for ten moles of NO produced the ΔH is 1,800 kJ. False; Easy

206.

CO2(g) + 2H2O(l) → CH4(g) + 2O2(g). This is an example of a formation reaction in the standard state.

False; Easy 207.

Na + ½ Cl2 → NaCl. This equation is an example of a formation reaction in the standard state.

True; Easy 208.

Given the thermochemical equation.

N2+ 2O2 → 2NO2 ∆H = –114.1 kJ The enthalpy of formation of NO 2 is –114.1 kJ. False; Easy 209.

True; Easy

The enthalpy of formation of O2 is 0 kJ.

210.

The enthalpy change of any chemical reaction is equal to the sum of the enthalpies of formation of the products minus the sum of the enthalpies of formation of the reactants.

True; Easy Multiple Choice Questions 211.

What is the SI unit of energy? a. joule b. watt c. ohm d. newton e. ampere a; Easy 212.

What is the kilojoule equivalent of 54 cal? a. 5.4  105 kJ b. 2.3  106 kJ c. 2.3  10-1 kJ d. 2.3  10-5 kJ e. 5.4  10-3 kJ c; Moderate 213.

What is the calorie equivalent of 562.0 J? a. 346.0 cal b. 134.0 cal c. 366.0 cal d. 155.0 cal e. 138.0 cal b; Easy 214.

Identify the joule equivalent of 4,184 calories. a. 22,020 J

b. 33,120 J c. 15,120 J d. 17,510 J e. 11,230 J d; Moderate 215.

What is the work performed by a gas, if it expands from 5.00 L to 15.0 L against a constant external pressure of 1.62 atm? Express the final answer in joules. a. –1,640 J b. –3,280 J c. –1,230 J d. –2,460 J e. –1,850 J a; Moderate 216.

What is the work performed when a gas expands from 56.00 mL to 83.00 mL against an external pressure of 760.0 torr? Express the final answer in joules. a. –10.52 J b. –20.52 J c. –3.900 J d. – 2.736 J e. –8.900 J d; Moderate 217.

The work done by a gas when it expands to 96 L is 4,100 J. If the external pressure is 2.0 atm, determine the initial volume of the gas. a. 56 L b. 16 L c. 20 L d. 36 L e. 76 L e; Moderate 218.

The work done by a gas when it expands from 90.0 L to 130.0 L is 2.03 kJ. What is the external pressure acting on the gas? a. 1.09  102 atm 1 b. 5.08  10 pa c. 3.10 torr d. 2.01  103 atm e. 2.04 pa b; Hard 219. _____ is the transfer of energy from one body to another due to a difference in temperature. a. Flux b. Enthalpy c. Magnetism d. Entropy e. Heat e; Easy

220.

In the equation, q = mcT, the constant ‗c‘ is called _____. a. entropy b. atomic radius c. specific heat capacity d. electronegativity e. enthalpy c; Easy 221.

What is the heat involved when 70.0 g of Ag increases its temperature from 90.0°C to 140.0°C? The specific heat of Ag is 0.233 J/g·°C. a. 583 J b. 105 J c. 1,050 J d. 816 J e. 1,630 J d; Easy 222.

The heat involved is 208.5 J when 30.0 g of Hg increases its temperature from _____ to 373°C. The specific heat of Hg is 0.139 J/g·°C. a. 353°C b. 175°C c. 50.0°C d. 315°C e. 323°C e; Easy 223.

It takes 460.0 J of heat to raise the temperature of 50.00 g of oxygen by 10.00°C. What is the specific heat of O2 in J/g·°C? a. 1.935 b. 2.690 c. 0.9200 d. 4.625 e. 1.916 c; Easy 224.

A mass of gold increases its temperature from 1.00 x102°C to 180.°C to produce 309.6 J heat. What mass of Au was present if the specific heat of Au is 0.129 J/g·°C. a. 10.0 g b. 15.0 g c. 30.0 g d. 45.0 g e. 60.0 g c; Easy 225.

The change in enthalpy equals heat at _____. a. absolute temperature b. constant pressure c. consistent volume d. constant length e. constant mass b; Easy

226. _____ refers to the heat of a process when pressure is held constant. a. Exothermic transformation b. Entropy change c. Specific heat d. Radiation e. Enthalpy change e; Easy 227.

A chemical equation that includes an enthalpy change is called a(n) _____. a. ionic equation b. hydro equation c. radioactive equation d. thermochemical equation e. electrochemical equation d; Easy 228.

Which of the following is a container used to measure the heat of a chemical reaction? a. manometer b. calorimeter c. barometer d. altimeter e. accelerometer b; Easy 229.

Calorimetry is the process of measuring _____ changes. a. enthalpy b. pressure c. density d. entropy e. volume a; Easy 230.

Which of the following thermochemical equations is used to represent the reaction of H2(g) with O2(g) to make H2O(ℓ) giving off 572 kJ? a. H2(g) + O2(g)  H2O(ℓ) ; H = −572 kJ b. H2(g) + O2(g)  2H2O(ℓ) ; H = +572 kJ c. H2(g) + O2(g)  2H2O(ℓ) ; H = −286 kJ d. 2H2(g) + O2(g)  2H2O(ℓ) ; H = −572 kJ e. 2H2(g) + 2O2(g)  4H2O(ℓ) ; H = −286 kJ d; Easy 231.

A solution of 0.100 mol of Ca2+ was mixed with 0.200 mol Cl– ions and CaCl2 was precipitated: Ca2+ + 2Cl–→ CaCl2 The temperature of the solution is increased by 21.00°C. What was the mass of the solution reacted if the enthalpy change for the production of 1 mol of CaCl2 is 176.0 kJ? (Assume that the solution has the same specific heat as water)? a. 11.00 g

b. 100.0 g c. 150.0 g d. 200.0 g e. 250.0 g d; Moderate 232.

To warm 800.0 g of H2O, 0.025 mol of phenol (C6H5OH) is burned. The water warms from 24.6C to 65.6C. What is the ΔH of the reaction on a molar basis? a. 4,110 kJ b. 5,480 kJ c. 1,600 kJ d. 2,740 kJ e. 1,370 kJ b; Moderate 233.

How much energy is given off or absorbed when 26.0 mol of O2 is reacted?

CH4(g) + 2O2(g) → CO2(g) + 2H2O(ℓ) ΔH = −890 kJ a. 28,900 kJ is absorbed b. 34,700 kJ is given off c. 5,790 kJ is absorbed d. 23,100 kJ is given off e. 11,600 kJ is given off e; Moderate 234.

Consider this thermochemical equation: CO2(g) + H2(g)  CO(g) + H2O(g) ΔH = 42 kJ How much energy is given off or absorbed when 22 g of CO2 is reacted? a. 42 kJ is absorbed b. 42 kJ is given off c. 21 kJ is given off d. 21 kJ is absorbed e. 22 kJ is absorbed d; Moderate 235.

Consider the thermochemical equation: N2 + O2 → 2NO; ΔH = +180.6 kJ How many moles of nitrogen should be reacted to absorb 6.00  102 kJ of energy? a. 2.77 mol b. 4.93 mol c. 3.32 mol d. 7.40 mol e. 14.8 mol c; Easy 236.

Consider the reaction: NH3 + HCl → NH4Cl ΔH = –176.0 kJ How much energy is given off when 170.3 g of NH4Cl is formed? (Molar mass of NH4Cl = 53.49 g/mol) a. 713.4 kJ b. 153.4 kJ c. 230.2 kJ

d. 560.3 kJ e. 1,274 kJ d; Easy 237.

Consider the thermochemical equation: 2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(ℓ); ΔH = –3,120 kJ How much energy is given off when 45.0 g of C2H6 and 96.0 g of oxygen are reacted?

(Molar masses: C2H6 = 30.0 g; O2 = 32.0 g) a. 1,340 kJ b. 3,120 kJ c. 1,560 kJ d. 782 kJ e. 2,670 kJ a; Hard 238.

Given the following thermochemical equation 2SO2 + O2→2SO3 ΔH = –395.2 kJ If 5,455 kJ is given off for the formation of SO3, what quantity of O2 in moles is reacted? a. 15.22 mol b. 4.180 mol c. 3.135 mol d. 13.80 mol e. 10.45 mol d; Easy 239.

Consider the thermochemical equation 2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(ℓ); ΔH = –3,120 kJ What amount of oxygen can consume 6,240 kJ of energy? (Molar masses: C2H6 = 30.0 g; O2 = 32.0 g)? a. 6.40  102 g b. 1.28  102 g c. 2.24  102 g d. 3.20  102 g e. 4.48  102 g e; Moderate 240.

Consider the thermochemical equation N2(g) + O2(g) → 2NO(g) ΔH = +180.6 kJ If 550 kJ of energy is supplied, determine the mass of nitrogen monoxide reacted. (Molar masses: N =14.00 g; O = 16.00 g) a. 14.0 g b. 28.0 g c. 84 g d. 182.7 g e. 180 g e; Moderate 241. _____ law states that when chemical equations are combined algebraically, their enthalpies can be combined in exactly the same way. a. Avogadro‘s

b. Henry‘s c. Charles‘s d. Boyle‘s e. Hess‘s e; Easy 242.

Given the thermochemical equation 2C(s) + 2O2(g) → 2CO2(g) ΔH = −787 kJ What is the ΔH for the reaction CO2(g) → C(s) + O2(g)? a. −0.00250 kJ b. −394 kJ c. +0.00250 kJ d. +394 kJ e. −787 kJ d; Easy 243.

Given the thermochemical equations, Ni + Cl2 → NiCl2 ΔH = −383 kJ NiCl2 + Cl2 → NiCl4 ΔH = −192 kJ what is the ΔH for 2NiCl2 → Ni + NiCl4? a. +191 kJ b. +575 kJ c. –191 kJ d. +383 kJ e. –575 kJ a; Easy 244.

Given the thermochemical equations, B2O3(s) + 3H2O(g) → 3O2(g) + B2H6(g) ΔH = −2,035 kJ H2O(ℓ) → H2O(g) ΔH = 44 kJ H2 (g) + ½O2(g) → H2O(ℓ) ΔH = −286 kJ 2B(s) + 3H2(g) → B2H6(g) ΔH = 36 kJ what is the ΔHf for 4B(s) + 3O2(g) → 2B2O3(s)? a. +1,273 kJ b. +2,690 kJ c. –1,273 kJ d. –5,594 kJ e. +636.5 kJ d; Hard 245.

For 2Na + Cl2 → 2NaCl ΔH is −772 kJ. What is the ΔH for the reaction NaCl → Na + ½Cl2? a. +386 kJ b. −772 kJ c. −386 kJ d. +772 kJ e. −579 kJ a; Easy 246.

If the ΔH for C3H6 + H2 → C3H8 is −125 kJ, what is the ΔH for the reaction 2C3H8 →2C3H6 + 2H2?

a. +250 kJ b. +125 kJ c. +62.5 kJ d. −125 kJ e. −250 kJ a; Easy 247.

Given the equations A + 6B → 4E ΔH = −234 kJ 2B +4C → 2E ΔH = −100 kJ What is the ΔH for the reaction A + 2B → 8C? a. +34 kJ b. −34 kJ c. +134 kJ d. +434 kJ e. −434 kJ b; Moderate 248. _____ reactions are chemical reactions that produce one mole of a substance from its constituent elements in their standard states. a. Formation b. Reduction c. Hydration d. Oxidation e. Combustion a; Easy 249.

What will be the coefficient of Cl2 when the following equation is written as a proper formation reaction?

H2 + Cl2 → 2HCl a. 1/4 b. 1/2 c. 1 d. 2 e. 4 b; Easy 250.

Which of the following is an example of a proper formation reaction? a. 2Na(s) + Cl2(g) → 2NaCl(s) b. 2H2(g) + O2(g) → 2H2O(ℓ) c. Ca(s) + C(s) + 3/2 O2(g) → CaCO3(s) d. 2Fe(s) + 3P(s) + 12O(g) → Fe2(PO4)3 e. H2(g) + Cl2(g) → 2HCl c; Easy 251.

Enthalpy change of any chemical reaction = _____ a. sum of the enthalpies of formation of the products + sum of the enthalpies of formation of the reactants b. 2(sum of the enthalpies of formation of the reactants) + 2(sum of the enthalpies of formation of the products)

c. sum of the enthalpies of formation of the products × sum of the enthalpies of formation of the reactants d. sum of the enthalpies of formation of the products − sum of the enthalpies of formation of the reactants e. sum of the enthalpies of formation of the products ÷ sum of the enthalpies of formation of the reactants d; Moderate 252.

Consider the equation.

N2(g)+ O2(g)→ 2NO(g) ∆H = 180.6 kJ What is the enthalpy of formation for NO? a. 90.3 kJ b. 66.2 kJ c. 99.3 kJ d. 49.4 kJ e. 132.4 kJ a; Easy 253.

The enthalpy of formation for C2H5OH is −277.0 kJ. What is the ∆H for the following reaction?

4C(s) + 6H2(g) + 2O2(g)→ 2C2H5OH(ℓ) a. +277.0 kJ b. −277.0 kJ c. −554.0 kJ d. +831.0 kJ e. +544.0 kJ c; Easy 254.

Identify the product that has zero enthalpy of formation. a. NO b. NO2 c. NH3 d. O2 e. O3 d; Easy 255.

Given the unbalanced reaction and the enthalpies of formation of molecules.

H2SO4 (ΔHf = −814.00 kJ/mol) + NaCl (ΔHf = −385.90 kJ/mol) → Na2SO4 (ΔHf = −331.64 kJ/mol) + HCl (ΔHf = −92.31 kJ/mol) What is the enthalpy change of the reaction? a. −2,670.80 kJ b. +1,069.54 kJ c. +2,670.80 kJ d. −852.35 kJ e. +1,898.49 kJ b; Hard Essay 256.

What is energy? What are its units?

Energy is the ability to do work. Work is defined as a force operating over a distance. In SI, force has units of newtons (N), while distance has units of meters. Therefore, work has units of N·m. This compound unit is redefined as a joule.

1 joule = 1 newton·meter 1 J = 1 N·m Calorie (cal) is also a common unit.

1 cal = 4.184 J Easy 257.

Describe systems and isolated systems. Provide examples.

A system is defined as the part of the universe under study. A beaker, a flask, or containers whose contents are being observed and measured are examples of systems.

An isolated system is a system that does not allow a transfer of energy or matter into or out of the system. A good approximation of an isolated system is a closed, insulated thermos-type bottle. The fact that the thermos-type bottle is closed keeps matter from moving in or out, and the fact that it is insulated keeps energy from moving in or out.

Easy 258.

Explain the relationship between external pressure and the work done by gases when expanding.

When a certain volume of a gas expands, it works against an external pressure to expand.

That is, the gas must perform work. Assuming that the external pressure Pext is constant, the amount of work done by the gas is given by the equation w = −Pext × ΔV Easy 259.

What is heat? How is it related to the temperature change and the mass of the object? What instrument is used to measure heat?

Heat is the transfer of energy from one body to another due to a difference in temperature. For a given object, the amount of heat (q) involved is proportional to two things: the mass of the object (m) and the temperature change (ΔT) evoked by the energy transfer. We can write this mathematically as q µ m ´ DT . The calorimeter is used to measure heat.

Easy 260.

What is specific heat capacity The amount of heat (q) involved is proportional to the mass of the object (m) and the temperature change (ΔT) evoked by the energy transfer. It can be written as q µ m ´ DT .

This proportionality can be made equality by multiplying with a constant as shown below. where  means ―is proportional to.‖ To make proportionality an equality, we include a proportionality constant. In this case, the proportionality constant is labeled c and is called the specific heat capacity, or, more succinctly, specific heat: q = mcΔT Easy 261.

What is enthalpy change? What is a thermochemical equation?

Enthalpy change is defined as the heat of a process when pressure is held constant.

A chemical equation that includes an enthalpy change is called a thermochemical equation.

Easy 262.

What are exothermic and endothermic reactions?

A chemical reaction that has a positive ΔH is an endothermic reaction, while a chemical reaction that has a negative ΔH is an exothermic reaction.

Easy 263.

How are ΔH values measured experimentally? ΔH is not measured; q is measured. But the measurements are performed under conditions of constant pressure, so ΔH is equal to the q measured.

Experimentally, q is measured by taking advantage of the equation q = mcΔT We premeasure the mass of the chemicals in a system. Then we let the chemical reaction occur and measure the change in temperature (ΔT) of the system. If we know the specific heat of the materials in the system (typically, we do), we can calculate q. That value of q is numerically equal to the ΔH of the process, which we can scale up to a molar scale.

The container in which the system resides is typically insulated, so any energy change goes into changing the temperature of the system, rather than being leaked from the system. The container is referred to as a calorimeter.

Moderate 264.

Write the equivalences of the equation H2(g) + Cl2(g) → 2HCl(g) ΔH = −184.6 kJ. 1 mol H2  1 mol Cl2 2 mol HCl  −184.6 kJ Easy 265.

Explain the concept of Hess‘s law.

Hess‘s law states that when chemical equations are combined algebraically, their enthalpies can be combined in exactly the same way. Two corollaries immediately present themselves: (1) If a chemical reaction is reversed, the sign on ΔH is changed.

(2) If a multiple of a chemical reaction is taken, the same multiple of the ΔH is taken as well.

Easy 266.

Consider the chemical equation C + O2 → CO2 ΔH = −393.5 kJ. How will the ΔH of the reaction change when the reaction is reversed? Explain.

If a chemical reaction is reversed, the sign on ΔH is changed.

Here the reaction becomes CO2 → C + O2. The ΔH of the reaction becomes + 393.5 kJ. Easy 267.

How would ΔH of a reaction change if the multiples of a chemical reaction are changed? Explain with an example.

If a multiple of a chemical reaction is taken, the same multiple of the ΔH is taken as well.

Consider the reaction Fe2(SO4)3 → Fe2O3 + 3SO3. The ΔH of the reaction is 570 kJ. If this equation is multiplied by five the ΔH should also be multiplied by five. The equation 5Fe2(SO4)3 → 5Fe2O3 + 15SO3 will have ΔH equal to 2,850 kJ. Easy 268.

What are formation reactions? What is a standard formation reaction?

Formation reactions are chemical reactions that form one mole of a substance from its constituent elements in their standard states. By standard states we mean as a diatomic molecule if that is how the element exists and the proper phase at normal temperatures.

Easy 269.

What is enthalpy of formation? Give an example.

The enthalpy change for a formation reaction is called the enthalpy of formation and is given the symbol ΔHf. Easy 270.

How do you determine the enthalpy change of a reaction from the enthalpies of formation of its components?

The enthalpy change of any chemical reaction is equal to the sum of the enthalpies of formation of the products minus the sum of the enthalpies of formation of the reactants.

DH rxn = å np DH f,p - å nr DHf,r

Easy Fill in the Blanks 271.

The ability to do work is called _____. energy; Easy 272.

40.84J = _____ cal.

9.761; Easy 273.

The statement that the total energy of an isolated system does not change is called the _____. law of conservation of energy; Easy 274.

The work done by a gas when expanding or contracting is given by the equation _____. w = −Pext × ΔV; Easy 275.

1 L·atm = _____ joules.

101.32; Easy 276. _____ is the transfer of energy from one body to another due to a difference in temperature.

Heat; Easy 277. _____ is defined as the heat of a process when pressure is held constant.

Enthalpy change; Easy 278.

A chemical equation that includes an enthalpy change is called a(n) _____. thermochemical equation; Easy 279.

The process of measuring changes in enthalpy is called _____. calorimetry; Easy 280.

Given:

H2(g) + Cl2(g) → 2HCl(g) ΔH = −184.6 kJ For 2.5 mole of Cl2 reacted, the enthalpy change is _____. −460 kJ; Easy 281.

Given: N2(g) + 3H2(g) → 2NH3(g) ΔH = −91.8 kJ 214.2 kJ energy is generated when _____ mol of H2 react.

7.000; Easy 282.

The ΔH of a reaction is 95 kJ. The ΔH will become _____ if the chemical reaction is reversed. −95 kJ; Easy 283.

If the ΔH for the reaction A + 2B → C is 150 kJ, ΔH for 6A + 12B → 6C is _____.

9.0x102 kJ; Easy 284.

The enthalpy change for a formation reaction is called _____. enthalpy of formation; Easy 285. _____ are chemical reactions that form one mole of a substance from its constituent elements in their standard states.

Formation reactions; Easy Chapter 8 Electronic Structure True/False Questions 286.

The wavelength of light is the number of cycles of light that pass a given point in one second.

False; Easy 287. one minute .

False; Easy

Frequency is the number of cycles of light that pass a given point in

288.

False; Easy

The frequency of light is 1016 s –1 if its wavelength is 10−8 m.

289.

True; Easy

The wavelength of light is 3 × 10–7 m if its frequency is 1015 s –1.

290.

True; Easy

The energy of light is 2.1 × 10−23 J if its frequency is 3.2 × 1010 s−1.

291.

A line spectrum can be obtained when the visible portion of the electromagnetic spectrum is passed through a prism.

False; Easy 292.

True; Easy

Frequency is represented by the Greek symbol, ν (Nu).

293.

The angular momentum quantum number affects the spatial distribution of the electron in three-dimensional space.

True; Easy 294. quantum number.

False; Easy

Electrons within a shell will have the same value of magnetic

295. and 1.

False; Easy

The spin quantum number of an electron should be between zero

296. numbers.

True; Easy

No two electrons in an atom can have the same set of four quantum

297.

False; Easy

Helium has two electrons. Its electron configuration is 1s11p1.

298.

The electron configuration for Ca (20 electrons) is 1s22s22p63s23p64s2.

True; Easy 299.

The electron configuration for Se (38 electrons) is 1s22s22p63s23p64s23d104p64d2.

False; Easy 300.

True; Easy

The 5s subshell is filled before allotting electrons to the 4d subshell.

301.

True; Easy

The first two columns of the periodic table are labeled the s block.

302.

False; Easy

Calcium is an element found in the d block of periodic table.

303.

False; Easy

The element Na has eleven valence electrons.

304.

True; Easy

The valence shell electron configuration of the element K is 4s1

305.

False; Easy

The valence shell electron configuration of S is 3s23p5.

306. table.

False; Easy

Atomic radii decrease as you go down a column of the periodic

307. decrease.

True; Easy

Going across a row on the periodic table, left to right, atomic radii

308.

False; Easy

The atomic radius of carbon is greater than that of boron.

309. left to right.

True; Easy

Ionization energies increase as you go across the periodic table from

310.

Electron affinity refers to the energy change when a gas-phase atom accepts an electron.

True; Easy Multiple Choice Questions 311.

The distance between corresponding points in two adjacent light cycles is called _____. k. quantum l. spectral line m. frequency n. spectrum o. wavelength e; Easy 312.

The number of cycles of light that pass a given point in one second is the_____ of light. f. wavelength g. spectrum h. spectral line i. frequency j. altitude d; Easy 313.

What is the frequency of light if its wavelength is 5.36 × 10−6 m? a. 2.12 × 102 s–1 b. 5.60 × 1013 s–1 c. 2.12 × 1014 s–1 d. 1.61 × 103 s–1 e. 1.79 × 10-14 s–1 b; Moderate 314.

What is the wavelength of light if its frequency is 1.6 × 1014 s–1? a. 5.3 × 10−7 m b. 1.9 × 10–6 m c. 4.8 × 1022 m d. 4.8 × 10–8 m e. 1.5 × 107 m b; Moderate

315.

What is the frequency of light if its wavelength is 813 nm? a. 3.69 × 1014 s–1 b. 2.71 × 10-14 s–1 c. 2.41 × 102 s–1 d. 8.10 × 1014 s–1 e. 2.71 × 10-15 s–1 a; Easy 316.

What is the energy of light if its frequency is 2.69 × 1010 s−1? a. 3.7 × 10–14 J b. 9.01 ×1016 J c. 1.78 × 10–23 J d. 2.69 × 10–34 J e. 2.69 × 10–24 J c; Moderate 317.

The energy of light is estimated to be 6.56 × 10−21 J. The frequency of this light is _____ s–1. a. 1.02 × 10–13 b. 4.34 × 100 c. 9.90 × 1013 d. 1.02 × 10–13 e. 9.90 × 1012 e; Easy 318.

What is the visible wavelength range in the electromagnetic spectrum? a. 800 nm to 1200 nm b. 400 nm to 700 nm c. 200 nm to 300 nm d. 50 nm to 500 nm e. 20 nm to 200 nm b; Easy 319.

Which of the following spectra appear when electricity is passed through a gas and light is emitted and this light is passed though a prism? a. Continuous spectrum b. Integrated spectrum c. Line spectrum d. Visible spectrum e. Absorption spectrum c; Easy 320.

Rainbow is an example of a(n) _____ spectrum. a. absorption b. line c. mass d. continuous e. power d; Easy 321.

Quantities that can only have certain specific values are called _____ values.

a. revoked b. quantized c. normalized d. optimized e. integrated b; Easy 322.

Which of the following is the reason why electrons have quantized energy values? a. Electrons have high wavelength b. Electrons emit electromagnetic radiation c. Electrons have high frequency d. Electrons have negative charge e. Electrons exist in specific orbits e; Moderate 323.

Which of the following statements is consistent with the theory of quantum mechanics? a. Electrons do not have specific wavelengths b. Electrons are collected into groups and subgroups c. Electrons act as waves rather than particles d. Electrons are neutrally charged particles e. Electrons are not randomly distributed around a nucleus c; Moderate 324.

The _____ quantum number largely determines the energy of an electron. a. magnetic b. parallel c. spin d. principal e. angular momentum d; Easy 325.

Which of the following is the quantum number that has a minor effect on the energy of the electron and affects the spatial distribution of the electron in threedimensional space? a. parallel b. principal c. magnetic d. spin e. angular momentum e; Easy 326.

Which of the following numbers can be the value for angular momentum quantum number if the principal quantum is 1? a. –1 b. 0 c. 1 d. 2 e. 3 b; Easy

327.

Which of the following magnetic quantum numbers is not possible if the principal quantum number is 4? a. 5 b. 2 c. 1 d. 0 e. –1 a; Moderate 328.

Each value of _____ quantum number designates a certain orbital. a. spin b. angular c. magnetic d. parallel e. principal c; Easy 329.

Which of the following sets of quantum numbers is shared by electrons in the same orbital? a. {2, 1, 0, −1/2} and {2, 1, 1, +1/2} b. {2, 1, 0, +1/2} and {2, 1, 0, −1/2} c. {3, 1, 1 , −1/2} and {3, 1, −1, +1/2} d. {3, 2, 0, +1/2} and {3, 2, 1, −1/2} e. {1, 1, −1, −1/2} and {1, 1, 1, + 1/2} b; Moderate 330.

Which of the following set of quantum numbers {n, ℓ, mℓ, ms} is allowed? a. {2, 1, 0, −1/2} b. {2, 1, 2, +1/2} c. {2, 2, 0, −1/2} d. {2, −1, 0, +1/2} e. {2, 2, 1, −1/2} a; Easy 331.

An electron is characterized by the following quantum number values. n = 3; ℓ = 2; mℓ = 1. Which of the following can be the spin quantum number for this electron? a. 0 b. 1 c. 2 d. −1/2 e. +1/6 d; Easy 332. a. b. c. d. e.

Which of the following is the central idea of the Pauli exclusion principle?

Electrons should have equal principal and magnetic quantum numbers.

At least three electrons should exist in a subshell.

Electrons have negative charge and positive magnetic polarization.

Electrons contain all colors of light.

No two electrons in an atom can have the same set of four quantum numbers.

e; Moderate 333.

Which of the following is a highest principal quantum number for the electrons in a calcium atom? a. 0 b. 1 c. 2 d. 3 e. 4 e; Easy 334.

How many subshells are completely filled with an atom that has 40 electrons? a. 9 b. 12 c. 8 d. 6 e. 4 a; Moderate 335.

How many subshells will have electrons if the atom has 15 electrons? a. 1 b. 2 c. 3 d. 4 e. 5 e; Easy 336.

Which of the following is the correct electronic configuration for an Mn atom with 25 electrons? a. 1s 2 2s 2 2 p 6 3s 2 3 p 6 4s 2 3d 5 b. 1s 2 2s 2 2 p 6 3s 2 3 p 6 3d 7 c. 1s 2 2s 2 2 p6 3s 2 3 p 6 4s 2 5s 2 3d 3 d. 1s 2 2s 2 2 p 6 3s 2 3 p 6 4s 2 3 p5 e. 1s 2 2s 2 2 p 6 3s 2 3 p 6 4s 2 4 p 4 3d 1 a; Moderate 337.

Identify the subshell that is expected to remain partially filled for an atom with 55 electrons. a. 4f b. 5p c. 4d d. 6s e. 5s d; Moderate 338.

The first two columns of the periodic table are labeled the _____. a. p block b. d block c. a block

d. f block e. s block e; Easy 339.

Which of the following elements belongs to the s block in the periodic table? a. Fe b. Ne c. C d. Na e. N d; Easy 340.

Which of the following refers to the p block in the periodic table? a. First two rows of the periodic table b. Right-most six columns of the periodic table c. First two columns of the periodic table d. Middle 10 columns of the periodic table e. The section that is detached from the main body of the table b; Easy 341.

Which of the following elements belong to the p block in the periodic table? a. Li b. Na c. O d. K e. H c; Moderate 342.

Electrons in the last unfilled subshell of an atom are called _____ electrons. a. inactive b. noble c. kinetic d. valence e. dormant d; Easy 343.

Which of the following is an element in the d block of the periodic table? a. Cl b. Fe c. Na d. K e. Ne b; Easy 344.

What are the valence shell electron configurations of the elements in the first column of the periodic table? a. np2 b. ns2 c. np1 d. np3 e. ns1

e; Easy 345.

How many valence electrons are present in a Ni atom with 28 electrons? a. 1 b. 2 c. 3 d. 5 e. 9 b; Easy 346.

Which of the following is the partially filled subshell for an atom that has 46 electrons? a. 4d8 b. 3d10 c. 4d10 d. 4f1 e. 5s2 a; Moderate 347.

What is the abbreviated electron configuration for S, which has 16 electrons? a. [Ar]4s1 b. [Ar]3p2 c. [Ne] 3s23p4 d. [Ne][Ar]4s2 e. [Ne]3p3 c; Moderate 348.

The inner electrons of an atom are called _____. a. noble electrons b. idle electrons c. redundant electrons d. core electrons e. valence electrons d; Easy 349.

Referring to the periodic table, which of the following atoms has the smallest atomic radius? a. Li b. Rb c. Na d. H e. K d; Easy 350.

Referring to the periodic table, which of the following atoms is likely to have the largest atomic radius? a. C b. O c. B d. F e. N

c; Easy 351.

Referring to the periodic table, which of the following atoms is likely to have the smallest atomic radius? a. Sc b. Cr c. Co d. Cu e. Zn e; Moderate 352. _____ is the amount of energy required to remove an electron from an atom in the gas phase. a. Electron affinity b. Kinetic energy c. Nuclear affinity d. Magnetic energy e. Ionization energy e; Easy 353.

Referring to the periodic table, which of the following atoms is likely to have the highest ionization energy? a. P b. S c. Al d. Cl e. Si d; Easy 354.

Referring to the periodic table, which of the following atoms is likely to have the highest ionization energy? a. F b. Cl c. Br d. I e. At a; Easy 355.

Predict which of the following atoms will have the highest magnitude of electron affinity. a. B b. F c. O d. N e. C b; Moderate Essay Questions 356.

Explain the concepts of wavelength and frequency of light. Draw a wave and show wavelength and frequency.

The wavelength of light is the distance between corresponding points in two adjacent light cycles, and the frequency of light is the number of cycles of light that pass a given point in one second. Wavelength is typically represented by λ, the lowercase Greek letter lambda, while frequency is represented by ν, the lowercase Greek letter nu.

Picture is upside down, can you fix this? I cannot.

Easy 357.

Explain the relationship between wavelength and frequency of light.

One property of waves is that their speed is equal to their wavelength times their frequency. That means we have speed = λν.

For light, speed is a universal constant when light is traveling through a vacuum. The measured speed of light (c) in a vacuum is about 3.00 × 108 m/s. Thus, we have c = λν Because the speed of light is a constant, the wavelength and the frequency of light are related to each other: as one increases, the other decreases and vice versa.

Easy 358.

What is electromagnetic spectrum? Which part of electromagnetic spectrum is visible?

Wavelengths, frequencies, and energies of light span a wide range; the entire range of possible values for light is called the electromagnetic spectrum. Light having a wavelength range between about 400 nm and 700 nm is visible. Light can have much longer and much shorter wavelengths than this, with corresponding variations in frequency and energy.

Easy 359.

Briefly distinguish between continuous spectrum and line spectrum.

An image that contains all colors of light is called continuous spectrum. A rainbow is an example of a continuous spectrum.

An image that contains only certain colors of light is called a line spectrum. It turns out that every element has its own unique, characteristic line spectrum.

Easy 360.

What is quantum number? Give a set of quantum numbers that exist.

Danish scientist Niels Bohr suggested that the electron in a hydrogen atom could not have any random energy, having only certain fixed values of energy that were indexed by the number n. The number n is called quantum number. Quantum number is an index that corresponds to a property of an electron, like its energy. {2,0,0,+1/2} Easy 361.

List the four types of quantum numbers.

The four types of quantum numbers are principal quantum number, angular momentum quantum number, magnetic quantum number, and spin quantum number.

Easy 362.

What is spin quantum number? What are the values that it can have?

Spin quantum number is the index that indicates one of two spin states for an electron.

Electrons and other subatomic particles behave as if they are spinning. Electrons themselves have two possible spin states, and because of mathematics, they are assigned the quantum numbers +1/2 and −1/2. These are the only two possible choices for the spin quantum number of an electron.

Easy 363.

Briefly explain Pauli‘s exclusion principle.

Pauli‘s exclusion principle says no two electrons in an atom can have the same set of four quantum numbers. This dramatically limits the number of electrons that can exist in a shell or a subshell.

Easy 364.

How would the eight electrons for O be assigned to the n and ℓ quantum numbers? Briefly explain.

The first two electrons go into the 1s shell-subshell combination. Two additional electrons can go into the 2s shell-subshell. The n = 2 shell also has a p subshell, so the remaining two electrons can go into the 2p subshell. The 2p subshell is not completely filled because it can hold a maximum of six electrons.

Easy 365.

What are s block elements? Provide an example.

The first two columns of the periodic table are labeled the s block. Hydrogen, lithium, sodium etc. are examples of s block elements.

Easy 366.

What are f block elements? Where does it appear in a periodic table?

The f block refers to the columns of the periodic table in which f subshells are being occupied. It is the 14-column section that is normally depicted as detached from the main body of the periodic table.

Easy 367.

Explain the concept of valence electrons and valence shell.

The electrons in the highest-numbered shell, plus any electrons in the last unfilled subshell, are called valence electrons; the highest-numbered shell is called the valence shell. The valence electrons largely control the chemistry of an atom. If we look at just the valence shell‘s electron configuration, we find that in each column, the valence shell‘s electron configuration is the same.

Easy 368.

What is atomic radius? Is it indicative of the size of an atom? How is it estimated?

Atoms behave as if they have a certain radius. The atomic radius is an indication of the size of an atom. Such radii can be estimated from various experimental techniques, such as the x-ray crystallography of crystals.

Easy

369.

What is ionization energy? How does it vary according to the relative position of elements in the periodic table?

Ionization energy (IE) is the amount of energy required to remove an electron from an atom in the gas phase. As you go down the periodic table, IE increases because the valence electron is farther away from the nucleus. As you go across the periodic table and the electrons get drawn closer in, it takes more energy to remove an electron; as a result, IE increases.

Easy 370.

What is referred to as electron affinity? How does it vary according to the relative position of elements in the periodic table?

Electron affinity is the energy change when a gas-phase atom accepts an electron. As you go across the periodic table, EA increases its magnitude. There is not a definitive trend as you go down the periodic table; sometimes EA increases, sometimes it decreases.

Easy Fill in the Blanks 371.

The distance between corresponding points in two adjacent light cycles is called _____. wavelength; Easy 372.

The measured speed of light in a vacuum is _________. .

2.9979 x 108m/s or 3.00 x 108 m/s; Easy 373.

The _____ refers to the full span of the possible wavelengths, frequencies, and energies of light. electromagnetic spectrum; Easy 374.

A _____ spectrum is an image that contains all colors of light. continuous; Easy 375.

A _____ is an index that corresponds to a property of an electron, like its energy. quantum number; Easy 376. _____ is a term used to describe electrons with the same principal quantum number.

Shell; Easy 377.

Each value of magnetic quantum number designates a certain _____. orbital; Easy 378.

The _____ principle states that no two electrons in an atom can have the same set of four quantum numbers.

Pauli exclusion; Easy 379.

A(n) _____ is the representation of the organization of electrons in shells and subshells in an atom. electron configuration; Easy

380.

The columns of the periodic table in which p subshells are being occupied are collectively called the _____. p block; Easy 381.

The electrons in the highest-numbered shell, plus any electrons in the last unfilled subshell are called _____. valence electrons; Easy 382.

The highest-numbered shell in an atom that contains electrons is known as the _____. valence shell; Easy 383. _____ refer to the variation of properties versus position on the periodic table.

Periodic trends; Easy 384.

The _____ is the amount of energy required to remove an electron from an atom in the gas phase. ionization energy; Easy 385. The _____ is the energy change when a gas-phase atom accepts an electron. electron affinity; Easy Chapter 9 Chemical Bonds True/False Questions 386. atom.

False; Easy

A Lewis electron dot diagram represents only the protons of an

387. atom.

The following is the accurate Lewis electron diagram for a carbon

False; Easy 388.

The Lewis electron diagram of an atom is shown below. This atom has a valence shell configuration of ns2np1.

True; Easy 389.

The following is the accurate Lewis electron diagram for a P atom.

False; Easy 390.

True; Easy

A Cl– ion has fulfilled the octet rule of valence electrons.

391.

A Na atom needs to gain two electrons to fulfill the octet rule.

False; Easy 392.

False; Easy

An ionic bond is formed between two atoms with the same charge.

393.

True; Easy

The measured strength of ionic bonding is called the lattice energy.

394.

Strength of an ionic bond decreases with an increase in the magnitude of the charge.

False; Easy 395.

When electrons are transferred to a central atom to form a compound, a covalent bond is formed.

False; Easy 396.

True; Easy

The H−H bond is an example of a covalent bond.

397.

True; Easy

O is the central atom in an H2O molecule.

398.

True; Easy

Lone electron pairs are electrons that do not make covalent bonds.

399.

The equal sharing of electrons in a covalent bond is called a nonpolar covalent bond.

True; Easy 400.

True; Easy

Electronegativity is a unitless number

401.

A bond will definitely be polar covalent if the electronegativity difference between the atoms involved is 1.4.

True; Easy 402.

False; Easy

A C–H bond is a nonpolar covalent bond.

403.

True; Easy

NO2 is an example of an odd electron molecule.

404.

False; Easy

Most odd electron compounds are chemically inactive.

405.

True; Easy

BeCl2 is an example of an electron deficient molecule.

406.

False; Easy

CH3OH is an example of a molecule that violates the octet rule.

407.

True; Easy

Any molecule with only two atoms has a linear shape.

408.

False; Easy

A GeF2 molecule has linear geometry.

409.

False; Easy

The molecule CH2O takes a trigonal pyramidal shape.

410. shape.

False; Easy

A molecule with four surrounding atoms takes a bent molecular

Multiple Choice Questions 411.

Which of the following is the common convention to represent a helium atom using the Lewis electron dot diagram? p. q. r. s. t. c; Easy 412.

A(n) _____ is a representation of the valence electrons of an atom that uses dots around the symbol of the element. k. Lewis electron dot diagram l. molecular electron graph m. structural electron formula n. molecular electron geometry o. orbital electron diagram a; Easy 413.

Which of the following is used to represent the Na atom using the Lewis electron dot diagram? a.

Na b.

Na c.

Na d.

Na e.

Na e; Easy 414.

Atoms like to have _____ electrons in their valence shell to remain stable. a. 8 b. 4 c. 3 d. 5 e. 6 a; Easy

415.

is the Lewis electron diagram of _____.

a. O2+ b. O c. O+ d. O– e. O2– e; Moderate 416.

How many electrons should a Zn atom lose to satisfy the octet rule? a. 3 b. 4 c. 0 d. 1 e. 2 e; Easy 417.

How many electrons should a Ca atom lose to satisfy the octet rule? a. 1 b. 2 c. 3 d. 4 e. 8 b; Easy 418. a. b. c. d.

For a chlorine atom to complete an octet, it must _____. lose an electron gain two electrons gain an electron gain five electrons

e. lose two electrons c; Easy 419.

What column of the periodic table has Lewis electron dot diagrams that have four electrons in them? a. the column headed by oxygen b. the column headed by fluorine c. the column headed by carbon d. the column headed by nitrogen e. the column headed by boron c; Moderate 420.

Which of the following ions contain an octet of valence electrons? a. N2– b. O– c. Na2+ d. F– e. Ba+ d; Moderate 421.

Which of the following ions do not contain an octet of valence electrons? a. Na+ b. Rb2+ c. S2– d. Cl– e. Ba2+ b; Moderate 422.

NaCl remains stable due to the attraction between the Na+ and Cl– ions. This attraction is called a(n)_____. a. chemical bond b. ionic bond c. covalent bond d. metallic bond e. dipole attraction b; Easy 423.

Which of the following occurs when Ca reacts with Cl2 to form CaCl2? a. Ca gains two electrons to become Ca2+ b. Cl2 loses two electrons to become Cl22– c. Ca loses two electrons to become Ca2+ d. Cl2 gains two electrons to become Cl22+ e. Cl atoms loses 2 electrons each to become Cl2– c; Easy 424. a. b. c. d. e.

The measured strength of ionic bonding is called the _____ energy. kinetic valence covalent lattice latent

d; Easy 425.

Which of the following compounds is characterized by the strongest ionic bond? a. LiCl b. NaCl c. NaBr d. LiF e. MgO e; Easy 426.

Which of the following bonds is formed when electrons are shared between atoms? a. ionic bond b. covalent bond c. dipole interaction d. valence bond e. metallic bond b; Easy 427.

Two hydrogen atoms share electrons to form a stable compound. This type of a bond is called a(n)_____ bond. a. ionic b. hydrogen c. dipole d. covalent e. metallic d; Easy 428.

Which of the following molecules is formed by triple bonds alone? f. F2 g. N2 h. O2 i. O3 j. CO2 b; Easy 429.

In the molecule CO2, how many electrons are shared by C and the two O atoms? a. 4 b. 5 c. 8 d. 6 e. 10 c; Easy 430.

Which of the following molecules is characterized by a single covalent bond? a. N2 b. F2 c. CO2 d. NaCl e. O2 b; Moderate

431.

is the Lewis electron diagram for _____.

f. BF4 g. BF4– h. BF42– i. BF4+ j. BF42+ b; Easy 432. f.

Identify the correct Lewis electron diagram for CO2.

g.

h.

i.

j.

d; Moderate 433. _____ is a scale for judging how much atoms of any element attract electrons. f. Electron density g. Reactivity h. Electron affinity i. Magnetivity j. Electronegetavity e; Easy 434.

If the electronegativity of a bond formed between two atoms is 0.15, it is likely to be a(n)_____ bond. a. likely ionic b. slightly polar covalent c. definitely polar covalent d. slightly ionic e. nonpolar covalent b; Easy

435.

Which of the following bonds is most likely to be nonpolar covalent bond? f. P–H g. C–H h. O–H (electronegativity chart in text shows a ―D‖ instead of ―O‖ for oxygen i. Na–Cl j. Li–Cl a; Easy 436.

What is the polarity of a C–Si bond? a. likely ionic b. slightly polar covalent c. definitely polar covalent d. slightly ionic e. nonpolar covalent c; Easy 437.

What is the polarity of a Na–Br bond? a. likely ionic b. slightly polar covalent c. definitely polar covalent d. slightly ionic e. nonpolar covalent c; Easy 438.

Which of the following bonds is most likely ionic? f. C–C g. Li–F h. P–H i. H–I j. H–Br b; Easy 439.

Which of the following has the strongest covalent bond? f. HF g. H2S h. H2 i. N2 j. CO2 d; Easy 440.

What is the energy change for this reaction? (Refer to Table 9.2 in the text.)

H

H

H

C

C

C

H

f. g. h. i.

H

–183 kJ/mol –477 kJ/mol 477 kJ/mol –129 kJ/mol

H

+

H

H

H

H

H

H

C

C

C

H

H

H

H

j. –4.00 kJ/mol a; Moderate 441.

What is the energy change when one carbon–carbon double bond is broken to form a carbon–carbon single bond and two H–H single bonds? (Refer to Table 9.2 in the text.) f. –565 kJ/mol g. –129 kJ/mol h. 609 kJ/mol i. –1,176 kJ/mol j. 129 kJ/mol c; Moderate 442.

Which of the following is a stable odd-electron molecule? a. N3 b. HF c. CO2 d. HCl e. NO e; Moderate 443.

Covalent compounds of _____ usually form electron deficient molecules. a. hydrogen b. boron c. chlorine d. lithium e. sodium b; Easy 444.

Which of the following compounds is characterized by electron deficient molecules? a. NaCl b. NO2 c. NO d. CO2 e. BeF2 e; Easy 445.

A covalent bond of any type is called a(n)_____. a. neutron group b. proton group c. core group d. nucleus group e. electron group e; Easy 446. a. b. c. d.

Which of the following molecules is not linear in shape?

PCl3 O2 CO2 NO

e. BeH2 a; Moderate 447.

Diatomic molecules are linear because _____. a. they form odd-electron molecules around the central atom b. electron-deficient molecules cannot be formed of only two electron groups c. only polar covalent bonds are formed in such molecules d. electron groups repel to get as far away from each other e. bonds in such compounds have a polarity of less than 0.4 d; Easy 448.

Which of the following molecules is likely to have a trigonal planar electron group distribution? a. H2O b. CO2 c. NO d. NO2 e. BF3 e; Easy 449.

From the following identify the molecule which has a bent or angular shape. a. CO2 b. CaCl2 c. GeF2 d. CaF2 e. SiO2 c; Easy 450.

What is the most likely shape for a molecule with a central atom and only one atom surrounding it? a. Angular b. Linear c. Trigonal planar d. Trigonal pyramidal e. Tetrahedral b; Moderate 451.

A molecule has three electron groups on the central atom and has two surrounding atoms. What is the likely shape of the atom? a. Trigonal pyramidal b. Bent c. Trigonal planar d. Linear e. Tetrahedral b; Moderate 452. a. b. c. d.

Which of the following molecules will likely have a bent molecular shape?

NOF CH2O CH2Cl2 CH4

e. PCl3 a; Moderate 453.

Identify the molecule with a tetrahedral molecular structure. a. NOF b. CH2O c. NH3 d. CH4 e. PCl3 d; Moderate 454.

Which of the following molecules will have a trigonal pyramidal molecular structure? a. PCl3 b. NO2 c. CH4 d. CaCl2 e. NaCl a; Moderate 455.

How many electron groups can likely be found on the central atom of a molecule with a trigonal planar shape? a. 1 b. 2 c. 3 d. 4 e. 5 c; Easy Essay Questions 456.

What are Lewis electron dot diagrams?

A Lewis electron dot diagram is a representation of the valence electrons of an atom that uses dots around the symbol of the element. The number of dots equals the number of valence electrons in the atom. These dots are arranged to the right and left and above and below the symbol, with no more than two dots on a side.

Easy 457.

Draw the Lewis electron diagram of a boron atom. Explain the diagram.

The Lewis electron diagram of a boron atom is shown below.

Boron has a valence shell electron configuration of 2s22p1. The two s electrons are represented on one side and the lone electron on the p shell is represented on the other side.

Easy 458.

Draw the electron diagram for an atom with the valence shell electron configuration of ns2np3.

The two s electrons should be represented together on one side.

Easy 459.

What is the octet rule? Do atoms always follow the octet rule when forming compounds?

The trend that atoms like to have eight electrons in their valence shell is called the octet rule. When atoms form compounds, the octet rule is not always satisfied for all atoms at all times, but it is a very good rule of thumb for understanding the kinds of bonding arrangements that atoms can make.

Easy 460.

Explain why it is difficult to violate the octet rule.

It is not impossible to violate the octet rule. Consider sodium: in its elemental form, it has one valence electron and is stable. It is rather reactive, however, and does not require a lot of energy to remove that electron to make the Na+ ion. We could remove another electron by adding even more energy to the ion, to make the Na2+ ion. However, that requires much more energy than is normally available in chemical reactions, so sodium stops at a 1+ charge after losing a single electron.

Easy 461.

What are the factors that determine the strength of ionic bonds?

The strength of ionic bonding depends on two major characteristics: the magnitude of the charges and the size of the ion. The greater the magnitude of the charge, the stronger the ionic bond. The smaller the ion, the stronger the ionic bond (because a smaller ion size allows the ions to get closer together). The measured strength of ionic bonding is called the lattice energy.

Easy 462.

What is a covalent bond? What is an ionic bond? Give an example that demonstrates each type of bond.

When electrons are shared between two atoms, they make a bond called a covalent bond.

The atoms engaged in a covalent bond can satisfy the octet rule and attain stability.

Example: CO2. The attraction between oppositely charged ions is called an ionic bond.

Ionic bonds are caused by electrons transferring from one atom to another. Example: NaCl Sodium transfers one electron to Chlorine which has seven electrons.

Moderate 463.

What is a single bond? Provide an example.

When two atoms are sharing one pair of electrons, this covalent bond is called a single bond.

Consider fluorine. F atoms have seven electrons in their valence shell. These two atoms can share their unpaired electrons to make a covalent bond as shown below.

Easy 464.

Explain the steps for determining the Lewis electron dot diagram of a simple molecule.

The central atom is the atom in the center of the molecule, while the surrounding atoms are the atoms making bonds to the central atom. The central atom is usually written first

in the formula of the compound (H2O is the notable exception). After the central and surrounding atoms have been identified, follow these steps: (1) Count the total number of valence electrons. Add extra if the species has negative charges and remove some for every positive charge on the species.

(2) Write the central atom and surround it with the surrounding atoms.

(3) Put a pair of electrons between the central atom and each surrounding atom.

(4) Complete the octets around the surrounding atoms (except for H).

(5)Put remaining electrons, if any, around the central atom.

(6) Check that every atom has a full valence shell.

Moderate 465.

What are polar covalent bonds and nonpolar covalent bonds?

The equal sharing of electrons in a covalent bond is called a nonpolar covalent bond. A covalent bond between different atoms that attract the shared electrons by different amounts and cause an imbalance of electron distribution is called a polar covalent bond.

Easy 466.

Explain the concept of electronegativity.

Electronegativity is a qualitative scale for judging how much atoms of any element attract electrons. Electronegativity is a unitless number; the higher the number, the more an atom attracts electrons.

Easy 467.

What are odd-electron molecules?

Some stable compounds have an odd number of electrons in their valence shells. Such molecules are called odd-electron molecules. Only a few of these kinds of compounds are stable. With an odd number of electrons, at least one atom in the molecule will have to violate the octet rule. Examples of stable odd-electron molecules are NO, NO2, and ClO2.

Easy 468.

What are electron-deficient molecules and expanded valence shell molecules?

Electron-deficient molecules have less than eight electrons around an atom in the molecule. The most common examples are the covalent compounds of beryllium and boron.

Some compounds have more than eight electrons assigned to their valence shell. These are called expanded valence shell molecules.

Easy 469.

Explain the concept of valence shell electron pair repulsion?

The basic idea in molecular shapes is called valence shell electron pair repulsion (VSEPR). It basically says that electron pairs, being composed of negatively charged particles, repel each other to get as far away from each other as possible. VSEPR makes a distinction between electron group geometry, which expresses how electron groups (bonds and nonbonding electron pairs) are arranged, and molecular geometry, which expresses how the atoms in a molecule are arranged.

Easy 470.

Briefly explain the trends in the formation of molecular shapes based on the number of electron groups and number of surrounding atoms.

The following table explains the general trend in the determination of molecular shapes.

Number of Electron Groups on Central Atom any 2 3 3 4 4 4

Number of Surrounding Atoms 1 2 3 2 4 3 2

Molecular Shape linear linear trigonal planar bent tetrahedral trigonal pyramidal bent

Easy Fill in the Blanks 471.

A _____ is a representation of the valence electrons of an atom that uses dots around the symbol of the element.

Lewis electron dot diagram; Easy 472.

The Lewis electron dot diagram of a Li atom is represented as _____. ; Easy

473.

The trend that atoms like to have eight electrons in their valence shell is called the _____. octet rule; Easy 474.

The attraction between oppositely charged ions is called a(n) _____. ionic bond; Easy 475.

The measured strength of ionic bonding is called the _____. lattice energy; Easy 476.

When electrons are shared between two atoms, they make a bond called a _____ bond. covalent; Easy 477.

When two atoms share one pair of electrons, the covalent bond is called a(n) _____. single bond; Easy 478.

The atoms making bonds to the central atom are called _____ atoms. surrounding atoms; Easy 479.

The unequal sharing of electrons in a covalent bond is called a _____ bond. polar covalent; Easy 480. _____ is a qualitative scale for judging how much atoms of any element attract electrons.

Electronegativity; Easy

481.

The approximate amount of energy needed to break a covalent bond is called the _____ of the covalent bond. bond energy; Easy 482.

In________, the number of electrons lost must equal the number of electrons gained. electron transfer; Easy 483.

A(n) _____ is a molecule with more than eight electrons in the valence shell of an atom. expanded valence shell molecule; Easy 484.

A CO2 molecule has a _____ shape. linear; Easy 485. A molecule has 4 electron groups on the central atom and has 3 surrounding atoms. The molecular shape of the molecule will be _____. trigonal pyramidal; Easy Chapter 10 Solids and Liquids True/False Questions 486. forces.

False; Easy

Dipole-dipole interactions are generally weaker than dispersion

487.

Polar molecules have a significantly higher boiling point than nonpolar molecules.

True; Easy 488.

True; Easy

H2O molecules experience hydrogen bonding.

489.

False; Easy

Hydrocarbons such as CH4 experience strong hydrogen bonding.

490. dispersion forces.

False; Easy

The strongest intermolecular forces present in a H2S molecule are

491.

True; Easy

Melting is always an endothermic reaction.

492.

True; Easy

ΔH will be negative for a solidification reaction.

493.

True; Easy

A substance stays at the same temperature during melting.

494.

True; Easy

The process of a liquid becoming a solid is called solidification.

495.

False; Easy

The process of a solid-to-gas change is called deposition.

496.

The energy change when 24 g of H2O(g) condenses to a liquid at 100°C is 40.68 kJ. (ΔHvap of H2O = 40.68 kJ/mol; molar mass of H2O = 18 g/mol) False; Easy 497.

Enthalpy of vaporization is the amount of energy needed to change from a liquid to a gas or from a gas to a liquid.

True; Easy 498.

The formation of a gas from a liquid at temperatures below the boiling point is called evaporation.

True; Easy 499.

The partial pressure of the vapor at equilibrium is called the vapor pressure of the liquid.

True; Easy 500. liquids.

False; Easy

Vapor pressures of solids are typically much higher than that of

501. temperature.

False; Easy

The vapor pressure for a substance is inversely proportional to the

502.

Surface tension is an effect caused by an imbalance of forces on the atoms at the surface of a liquid.

True; Easy 503.

False; Easy

Adhesion is the tendency of a substance to interact with itself.

504.

False; Easy

Mercury has stronger adhesive forces than cohesive forces.

505.

False; Easy

Solids can be easily compressed.

506.

True; Easy

An amorphous solid has no long-term structure or repetition.

507. structure.

True; Easy

A crystalline solid has a regular, repeating three-dimensional

508.

Ionic solids are conductors of electricity in their solid state.

False; Easy 509.

True; Easy

Carbon in its diamond form is a covalent network solid.

510.

True; Easy

The element Ag will exist as a metallic solid in the solid state.

Multiple Choice Questions 511.

The preferred phase of a substance at a given set of conditions is a balance between the _____ of the particles and intermolecular forces between the particles. p. magnetic properties q. energy r. flammability s. wavelength t. frequency b; Easy 512. _____ is caused by the instantaneous position of an electron in a molecule, which temporarily makes that point of the molecule negatively charged and the rest of the molecule positively charged. f. Dipole-dipole interaction g. Hydrogen bonding h. Ionic bonding i. Covalent bonding j. Dispersion force e; Easy 513.

Considering the dispersion force, which of the following molecules is least likely to be a gas? a. calcium b. neon c. nitrogen d. oxygen e. neon a; Easy 514.

By looking at the periodic table, identify the atom that would experience the strongest dispersion force. a. N b. O c. Al d. C e. Cl e; Easy 515.

Identify the substance where the most significant intermolecular force would be the dispersion force. k. HCl l. HBr

m. C2H6 n. NaOH o. NH3 c; Easy 516.

Which of the following substances experience hydrogen bonding? f. H2O g. HCl h. CH4 i. H2S j. C6H6 a; Easy 517.

Identify the substance for which the most significant intermolecular force will be the dispersion force. f. NaCl g. NH3 h. HCl i. C3H8 j. (CH3)2CO d; Easy 518.

For which of the following products will the most significant intermolecular force NOT be the dipole-dipole interaction k. HF l. C6H5OH m. HI n. NH3 o. CH4 e; Easy 519.

Hydrogen bonds are formed when hydrogen is bonded to a(n) _____ atom. k. sodium l. chlorine m. bromine n. oxygen o. iodine d; Easy 520.

The process of a solid becoming a liquid is called _____. k. boiling l. normalization m. solidification n. freezing o. melting e; Easy 521.

The temperature at which a substance in the solid state becomes a liquid is called the_____. a. heat of formation b. latent point

c. melting point d. boiling point e. entropy of fusion c; Easy 522.

Which of the following at best describes the enthalpy of fusion? f. Energy needed to change from a liquid to gas g. Temperature at which melting occurs h. Temperature at which boiling occurs i. Energy needed to change from a solid to liquid j. Temperature at which a solid becomes conductive d; Easy 523.

What is the energy change when 1.50 mol of H2O melts at 273 K? (ΔHfus of H2O = 6.01 kJ/mol) a. 9.02 kJ b. 6.01 kJ c. 3.01 kJ d. 12.0 kJ e. 4.52 kJ a; Easy 524.

24.875 kJ of energy change is observed when benzene melts at 5.5°C. How many grams of benzene are melted? (Molar mass of C6H6 =78 g/mol; ΔHfus = 9.95 kJ/mol) a. 36 g b. 72 g c. 108 g d. 144 g e. 195 g e; Easy 525.

A reaction where the substance remains at the same temperature is called a(n) _____. a. adiabatic process b. isothermal process c. isobaric process d. polytrophic process e. isochoric process b; Easy 526.

What will happen to liquids when stored in zero gravity environments? a. Liquids take the shape of the container. b. Liquids harden and become solid particles. c. Liquids float in such environments. d. Liquids expand to become a larger mass. e. Liquids evaporate to become gases. c; Easy 527.

Which of the following refers to the amount of energy required to convert a liquid to a gas? a. Enthalpy of formation

b. Heat of fusion c. Enthalpy of fusion d. Heat of liquidation e. Enthalpy of vaporization e; Easy 528.

What is the energy change when 72 g of H2O (g) condense to a liquid at 373 K? (ΔHvap of water = 40.68 kJ/mol; molar mass of H2O = 18 g/mol) a. 81 kJ b. 41 kJ c. 160 kJ d. 210 kJ e. 240 kJ c; Easy 529.

The energy change occurring is 2.10 kJ when_____ of H2O (g) condenses to a liquid at 100°C. (ΔHvap of water = 40.68 kJ/mol; molar mass of H2O = 18 g/mol). a. 12.6 g b. 8.56 g c. 1.63 g d. 17.6 g e. 0.929 g e; Easy 530.

What is the energy change when 565 g of C2H5OH boil at its normal boiling point of 78.3°C? (ΔHvap for ethanol = 38.6 kJ/mol; molar mass of C2H5OH = 46.0 g/mol) a. 396 kJ b. 302 kJ c. 119 kJ d. 474 kJ e. 124 kJ d; Easy 531.

Which of the following terms refer to the solid-to-gas change of compounds? a. Boiling b. Normalization c. Solidification d. Deposition e. Sublimation e; Easy 532.

Identify the term that refers to the gas-to-solid change of compounds. a. Deposition b. Normalization c. Solidification d. Sublimation e. Boiling a; Easy 533.

The ΔHvap for water is 40.68 kJ/mol and the ΔHfus is 6.010 kJ/mol. Calculate the ΔHsub.

a. 34.67 kJ/mol b. 46.69 kJ/mol c. 244.5 kJ/mol d. 69.34 kJ/mol e. 93.38 kJ/mol b; Easy 534.

The formation of a gas from a liquid at temperatures below the boiling point is called _____. a. sublimation b. deposition c. evaporation d. melting e. under-boiling c; Easy 535.

Which of the following is considered the reverse process of vaporization? a. osmosis b. solidification c. melting d. sublimation e. condensation e; Easy 536.

The temperature at which the vapor pressure of a liquid equals the surrounding environmental pressure is called the _____. a. enthalpy of fusion b. boiling point c. melting point d. enthalpy of vaporization e. freezing point b; Easy 537.

The normal vapor pressure is the temperature at which the vapor pressure is _____ torr. a. 460 b. 500 c. 600 d. 760 e. 800 d; Easy 538.

By referring to the plots of vapor pressure versus temperature provided in the text, identify the likely boiling point of water at 600 torr vapor pressure. a. 92°C b. 98°C c. 56°C d. 69°C e. 78°C a; Easy

539. _____ is responsible for the fact that small insects can ―walk‖ on water. a. Vapor pressure b. Surface tension c. Change in enthalpy d. Density e. Viscosity b; Moderate 540.

Certain liquids tend to bead up when present in small amounts. What is the reason for this? a. low surface tension b. high vapor pressure c. low vapor pressure d. high surface tension e. low density d; Easy 541.

Liquids form spheres in free fall or zero gravity. Which of the following is the reason for this behavior? a. liquid vapor pressure b. surface tension c. viscosity d. low density e. entropy b; Easy 542. _____ is the tendency of a substance to interact with other substances because of intermolecular forces. a. Adhesion b. Viscosity c. Osmosis d. Capillary action e. Cohesion a; Easy 543. _____ is the tendency of a substance to interact with itself. a. Surface tension b. Adhesion c. Osmosis d. Cohesion e. Viscosity d; Easy 544.

When a liquid is introduced to a small-diameter tube of another substance, the liquid moves up or down in the tube, as if ignoring gravity. Which of the following terms specifically refers to this action? a. adhesion b. capillary action c. surface tension d. cohesion e. viscosity

b; Easy 545.

Which of the following is a major similarity between solids and liquids? a. Strength of the intermolecular forces is the same in both the phases b. Covalent bonds are not common in both the phases c. The particles are in contact with each other in both the phases d. Hydrogen bonds are not common in both the phases e. They can take the shape of the container in which they are stored c; Easy 546.

A(n) _____ solid is a solid with no long-term structure or repetition. a. amorphous b. crystalline c. molecular d. ionic e. covalent network a; Easy 547.

A(n) _____ solid is a solid that has a regular, repeating three-dimensional structure. a. amorphous b. molecular c. ionic d. covalent network e. crystalline e; Easy 548.

Which of the following is an example of an amorphous solid? a. metals b. biological molecules c. glasses d. salt crystals e. semiconductors c; Easy 549.

An NaCl crystal is an example of a(n) _____ solid a. crystalline b. ionic c. molecular d. amorphous e. covalent network b; Easy 550.

Which of the following is an ionic solid? a. NaBr b. I2 c. C6H6 d. S5 e. C6Cl6 a; Easy

551.

Which of the following types of bonding is commonly seen in molecular solids? a. hydrogen bond b. ionic bond c. ion-dipole bond d. covalent bond e. dipole interaction d; Easy 552.

Which of the following is a molecular solid? a. CuCl b. AuCl c. NaCl d. CuCl2 e. CCl4 e; Easy 553.

What kind of crystal would CH4 form in its solid state? a. ionic b. molecular c. amorphous d. crystalline e. covalent network b; Easy 554.

Which of the following elements will form a metallic solid in its solid state? a. O b. S c. Se d. Cu e. C d; Easy 555.

A shiny and silvery solid substance stored in the form of a sheet is a(n) _____ solid. a. amorphous b. ionic c. metallic d. crystalline e. covalent network c; Easy Essay Questions 556.

Why does a substance have the phase it does? How does the temperature affect the existence in various phases?

The preferred phase of a substance at a given set of conditions is a balance between the energy of the particles and intermolecular forces (or intermolecular interactions) between the particles. If the forces between particles are strong enough, the substance is a liquid or, if stronger, a solid. If the forces between particles are weak and sufficient energy is present, the particles separate from each other, so the gas phase is the preferred phase.

The energy of the particles is mostly determined by temperature, so temperature is the main variable that determines what phase is stable at any given point.

Easy 557.

What is dispersion force? Explain.

Dispersion force is an intermolecular force caused by the instantaneous position of an electron in a molecule. This interaction is caused by the instantaneous position of an electron in a molecule, which temporarily makes that point of the molecule negatively charged and the rest of the molecule positively charged. In an instant, the electron is now somewhere else, but the fleeting imbalance of electric charge in the molecule allows molecules to interact with each other.

Easy 558.

What are dipole-dipole interactions?

Molecules with a permanent dipole moment experience dipole-dipole interactions, which are generally stronger than dispersion forces if all other things are equal. The oppositely charged ends of a polar molecule, which have partial charges on them, attract each other.

Easy 559.

Does hydrogen bonding exist in both H2S and NH3? Explain your answer.

Hydrogen bonding is a strong interaction between molecules due to H atoms being bonded to N, O, or F atoms. In NH3, hydrogen bonding is formed between hydrogen and nitrogen but hydrogen bonding is not formed between S and H. Easy 560.

Explain the terms melting, solidification, and melting point.

The process of a solid becoming a liquid is called melting (an older term that you may see sometimes is fusion). The opposite process, a liquid becoming a solid, is called solidification. For any pure substance, the temperature at which melting occurs is known as the melting point.

Easy 561.

What is enthalpy of fusion? Provide an example.

Enthalpy of fusion (or heat of fusion) is the amount of energy needed to change from a solid to a liquid or from a liquid to a solid. It is referred to as ΔHfus.

The ΔHfus for water is 6.01 kJ/mol.

Easy 562.

Explain how boiling point changes with changes in pressure.

The liquid/gas conversion process is noticeably affected by the surrounding pressure on the liquid because gases are strongly affected by pressure. This means that the temperature at which a liquid becomes a gas, the boiling point, can change with surrounding pressure. Therefore, we define the normal boiling point as the temperature at which a liquid changes to a gas when the surrounding pressure is exactly 1 atm, or 760 torr.

Easy 563.

What are sublimation and deposition? What is enthalpy of sublimation?

The solid-to-gas change is called sublimation, while the reverse process is called deposition. Sublimation is isothermal, like the other phase changes. There is a measurable

energy change during sublimation; this energy change is called the enthalpy of sublimation, represented as ΔHsub.

The relationship between the ΔHsub and the other enthalpy changes is as follows: ΔHsub = ΔHfus + ΔHvap Easy 564.

Explain the process of evaporation.

The formation of a gas from a liquid at temperatures below the boiling point is called evaporation. At these temperatures, the material in the gas phase is called vapor, rather than gas; the term gas is reserved for when the gas phase is the stable phase.

Easy 565.

What is surface tension? List some of its effects on liquids.

Surface tension is an effect caused by an imbalance of forces on the atoms at the surface of a liquid.

Liquids with high surface tension tend to bead up when present in small amounts. Surface tension causes liquids to form spheres in free fall or zero gravity. Surface tension is also responsible for the fact that small insects can ―walk‖ on water.

Easy 566.

What are adhesion and cohesion? What are their impacts?

Adhesion is the tendency of a substance to interact with other substances because of intermolecular forces, while cohesion is the tendency of a substance to interact with itself. If cohesive forces within a liquid are stronger than adhesive forces between a liquid and another substance, then the liquid tends to keep to itself; it will bead up. However, if adhesive forces between a liquid and another substance are stronger than cohesive forces, then the liquid will spread out over the other substance, trying to maximize the interface between the other substance and the liquid.

Easy 567.

Explain the concept of capillary action.

Capillary action is the behavior of a liquid in narrow spaces due to differences in adhesion and cohesion.

Easy 568.

What are amorphous solids? Provide examples.

An amorphous solid is a solid with no long-term structure or repetition. Examples include glass and many plastics, both of which are composed of long chains of molecules with no order from one molecule to the next.

Easy 569.

What are ionic solids? Explain with an example.

An ionic solid is a crystalline solid composed of ions (even if the ions are polyatomic).

NaCl is an example of an ionic solid. The Na+ ions and Cl− ions alternate in three dimensions, repeating a pattern that goes on throughout the sample. The ions are held together by the attraction of opposite charges—a very strong force. Hence most ionic solids have relatively high melting points.

Easy 570.

What are molecular solids? Explain their behavior. Provide an example.

A molecular solid is a crystalline solid whose components are covalently bonded molecules. Many molecular substances, especially when carefully solidified from the liquid state, form solids where the molecules line up with a regular fashion similar to an ionic crystal, but they are composed of molecules instead of ions. Because the intermolecular forces between molecules are typically less strong than in ionic solids, molecular solids typically melt at lower temperatures and are softer than ionic solids. Ice is an example of a molecular solid.

Easy Fill in the Blanks 571. _____ is an intermolecular force caused by the instantaneous position of an electron in a molecule.

Dispersion force; Easy 572. _____ is an intermolecular force caused by molecules with a permanent dipole.

Dipole-dipole interaction; Easy 573.

The process of a solid becoming a liquid is called _____. melting; Easy 574.

The characteristic temperature at which a solid becomes a liquid is called the _____. melting point; Easy 575. _____ refers to the amount of energy needed to change from a solid to a liquid or from a liquid to a solid.

Enthalpy of fusion; Easy 576.

The process of a gas becoming a liquid is called _____. condensation; Easy 577.

The amount of energy needed to change from a liquid to a gas or from a gas to a liquid is referred to as _____. enthalpy of vaporization; Easy 578.

The process of a gas becoming a solid is called _____. deposition; Easy 579.

The formation of a gas phase from a liquid at temperatures below the boiling point is called _____. evaporation; Easy 580. _____ refers to a situation in which a process still occurs but the opposite process also occurs at the same rate so that there is no net change in the system.

Dynamic equilibrium; Easy 581. _____ refers to an effect caused by an imbalance of forces on the atoms at the surface of a liquid.

Surface tension; Easy

582.

The tendency of a substance to interact with other substances because of intermolecular forces is called _____. adhesion; Easy 583.

A solid with a regular, repeating three-dimensional structure is known as a(n) _____ solid. crystalline; Easy 584. _____ are crystalline solids whose components are covalently bonded molecules.

Molecular solids; Easy 585. A solid with the characteristic properties of a metal is called a(n) _____ solid. metallic; Easy Chapter 11 Solutions True/False Questions 586.

Solutions exist in all phases of matter.

True; Easy

587.

CH3OH is not soluble in water False; Easy

588.

NaOH will be more soluble in CCl4 than H2O. False; Easy

589.

Molarity of a solution is the number of moles of solute divided by the milliliters of the solution.

False; Easy 590.

The molarity of a solution is 1.25 M when 50.0 g of NaOH are dissolved to make 1.00 L of the solution.

True; Moderate 591.

0.5 mole of a solute is present in 0.3 L of a 5 M NaCl solution.

False; Moderate

592. solvent.

False; Easy

Molality refers to the number of moles of solute per gram of

593.

For ionic solutions, the total ion concentration is the sum of the individual ion concentrations.

True; Easy 594. same.

True; Easy

In both dilution and concentration, the amount of solute stays the

595.

The molarity of a solution is 25 M if 5.0 moles of the solute is present in 5.0. liters of the solution.

False; Easy 596.

If 25.0 mL of a 1.0 M solution is diluted to 50.0 mL, the concentration becomes 0.50 M. True; Easy 597. reduces by 66.7%.

False; Easy

If 10 mL of solvent is added to 20 mL of solution, the molarity

598. molecules.

False; Easy

15.0 liters of 10.0 M solution of HCl contains 0.667 moles of HCl

599.

20.0 g NaOH is present in 2.0 liters of 0.25 molar solution of NaOH. (Molar mass of NaOH = 40 g/mol) True; Easy 600.

39.19 g NaCl is diluted to create a 1.000 M solution. The volume of the solution is 0.5845 L. True; Easy 601.

A solution is made by mixing 2.00 mol of solute and 15.0 mol of solvent. The mole fraction of the solute is 0.133 False; Easy 602.

Solutions have a lower vapor pressure than the pure solvent.

True; Easy

603.

Solutions have a higher boiling point than pure solvents.

True; Easy

604.

The boiling point of a 5.00 m solution of C6H4Cl2 in CCl4 is 101.6°C. Assume that C6H4Cl2 is not volatile and Kb = 4.95°C/m.

True; Easy 605.

Solutions have a higher freezing point than their pure solvents.

False; Easy

606.

0°C. True; Easy

The freezing point of a 0.5 m solution of NaCl is water is less than

607.

Osmotic pressure is the tendency of a solution to pass solution through a semipermeable membrane due to concentration differences.

False; Easy The van‘t Hoff factor for CCl4 is 5.

608.

False; Easy

The van‘t Hoff factor for CaCO3 is 2.

609.

True; Easy 610.

Kf = 1.86°C/m.

True; Easy

The freezing point of a 1.9 m solution of NaCl in H2O is –3.5°C if

Multiple Choice Questions 611. Which of the following terms refer to the major component of a solution? u. solute v. solvent w. precipitate x. stream y. pour out b; Easy 612. A concentrated solution refers to a solution that has _____. a. very low density b. very high viscosity c. large amounts of solute d. large amounts of solvent e. very high density c; Easy 613. Only a certain amount of solute can be dissolved in a given amount of solvent. This maximum amount is called the _____ of the solute. a. solubility b. reactivity c. dilution d. hydraulicity e. acidity a; Easy 614. Identify the solvent in which C2H5OH will be most soluble. a. CH4 b. CCl4 c. C6H6 d. H2O e. CO2 d; Easy 615. The number of moles of solute divided by the number of liters of solution is _____. a. concentration b. molar volume c. dilution d. volume deliberation e. molarity e; Easy

616. What is the molarity of a solution made when 46.76 g of NaCl is dissolved to make 500.0 mL of solution? (Molar mass of NaCl = 58.45 g/mol) a. 2.500 M b. 0.002500 M c. 1.600 M d. 80.00 M e. 0.001600 M c; Easy 617. The molarity of a solution made by dissolving 2.00 moles of HCl in water is 2.65 M. What is the volume of the solution? a. 2.65 L b. 755 mL c. 265 mL d. 0.53 L e. 53 mL b; Easy 618. How many moles of a solute are present in 200.00 mL of a 1.50 M HCl solution? a. 0.300 mol b. 3.00 mol c. 3.00 × 102 mol d. 75.0 mol e. 7.50 mol a; Easy 619. How much solute is present in 5.00 L of a 1.50 M solution? (Molar mass of solute = 98 g/mol) a. 735 g b. 749.0 g c. 708 g d. 558.8 g e. 440.8 g a; Easy 620. 80.0 g of a solute is present in 9.00 L of a 0.200 M solution. What is the molar mass of the solute? a. 80.0 g/mole b. 44.4 g/mole c. 1.78 g/mole d. 17.8 g/mole e. 73.6 g/mole b; Easy 621. 50.0 g of Fe is present in 1.50 kg of a solvent. What is the molality of the solution?

(Molar mass of Fe = 55.8 g/mol) a. 0.896 m b. 0.597 m c. 1.12 m d. 1.34 m e. 0.744 m

b; Easy 622. If there is 9.00 g of Ag present in 190 g of solution, what is the Ag concentration in parts per thousand? a. 21.0 ppth b. 1.70 ppth c. 47 ppth d. 47.4 ppth e. 17.0 ppth c; Easy 623. The concentration of Br– ion in a sample of H2O is 10 ppm. What mass of Br– ion is present in 300.0 mL of H2O, which has a density of 1.02 g/mL? a. 10 mg b. 6.00 mg c. 3.00 mg d. 3.06 mg e. 4.52 mg d; Easy 624. The concentration of Cl– ion in a sample of H2O is 21.0 ppth. What mass of Cl– ion is present in 0.900 kg of H2O? a. 9.45 g b. 9.45 mg c. 18.9 mg d. 21 g e. 18.9 g e; Easy 625. The process of adding solvent to a solution is called _____. a. concentration b. saturation c. precipitation d. ionization e. dilution e; Easy 626. If 3.0 mL of a 1.0 M solution is diluted to 12 mL, what is the final concentration? a. 4.0 M b. 0.25 M c. 1.0 M d. 0.50 M e. 2.0 M b; Easy 627. 50.0 mL of a 1.00 M solution is diluted to 0.300 M. What is the final volume of the solution? a. 150 mL b. 12.5 mL c. 167 mL d. 200 mL

e. 100 mL c; Easy 628. 6.00 L of a 0.400 M solution of HCl should be diluted to a 0.250 M solution. What amount of water should be added to get the desired concentration? a. 7.55 L b. 8.00 L c. 3.60 L d. 9.00 L e. 9.60 L c; Easy 629. 10.0 L of water is added to 15.0 L of a 1.00 M solution of H2SO4. What is the new molarity? a. 2.50 M b. 1.00 M c. 1.50 M d. 0.600 M e. 0.667 M c; Easy 630. 12.00 L of water is added to 18.00 L of solution to change its molarity to 1.000 M. What was the initial molarity? a. 1.667 M b. 1.500 M c. 1.465 M d. 2.161 M e. 3.334 M b; Easy 631. 10.0 L of HCl is diluted by mixing with water. What is the amount of water added if the molar concentration has become 1/3rd the initial molar concentration? a. 3.33 L b. 2.50 L c. 16.7 L d. 13.3 L e. 20.0 L e; Easy 632. 750 mL of water is added to a 750 mL solution to reduce its concentration. What is the percentage change in its molar concentration? a. 25.0% b. 33.3% c. 50.0% d. 66.7% e. 100% c; Easy 633. How many moles of HCl are present in 1.20 L of a 0.850 M solution? a. 2.05 mol b. 1.02 mol

c. 0.980 mol d. 0.708 mol e. 1.68 mol b; Easy 634. What is the molarity of a solution if 15 g NaCl is present in 0.50 L of the solution?

(Molar mass of NaCl = 58.45 g/mol) a. 0.26 M b. 7.5 M c. 15 M d. 0.51M e. 3.0 M d; Easy 635. What mass of CaCl2 is needed to produce 15.0 liters of 1.60 M solution of CaCl2? (Molar mass of CaCl2 = 111 g) a. 2.40 kg b. 2.66 kg c. 1,040 g d. 24.0 g e. 2.08 kg b; Easy 636. How much NaCl is present in 15 L of a 1.5 M solution of NaCl? (Molar mass of NaCl = 58.5 g/mol) a. 1.32 kg b. 22.5 kg c. 2.25 kg d. 315 g e. 565 g a; Easy 637. What mass of solute is present in 2.050 L of 0.6630 M NaOH? (Molar mass of NaOH = 40.00 g/mol) a. 82.00 g b. 54.37 g c. 66.30 g d. 13.26 g e. 10.25 g b; Easy 638. What volume of 0.40 M NaOH will react with 20.0 g of H2C2O4(s) according to the following chemical equation? (Molar mass of H2C2O4 is 90.0 g/mol) H2C2O4(s) + 2NaOH(aq) → Na2C2O4(aq) + 2H2O(ℓ) a. 0.422 L b. 0.211 L c. 1.11 L d. 0.169 L e. 0.339 L c; Easy

639. 1.70 L of 0.550 M H2O2 reacts according to the following chemical reaction.

2H2O2(aq) → 2H2O(ℓ) + O2(g) How much H2O is formed? (Molar mass of H = 1.00 g/mol; O = 16.0 g/mol) a. 0.935 kg b. 3.65 g c. 1.87 g d. 16.8 g e. 0.935 g d; Easy 640. _____ refers to a property of solutions related to the fraction that the solute particles occupy in the solution, not their identity. a. Dissolutive property b. Static property c. Dynamic property d. Kinetic property e. Colligative property e; Easy 641. A solution is made by dissolving 15.00 g of NaCl in 500.0 g of H2O. What is the mole fraction of NaCl in the solution? a. 0.0089 b. 0.2566 c. 0.0092 d. 0.1156 e. 0.2778 c; Easy 642. A solution contains 65.0 g of solvent. How much solute is present in the solution if the mole fraction of the solute is 0.135? (Molar mass of solvent = 18 g/mol; molar mass of solute = 30 g/mol) a. 0.563 g b. 4.05 g c. 3.61 g d. 53.3 g e. 16.9 g e; Moderate 643. A solution is made by mixing 25.0 g of NaOH in H2O. What is the amount of H2O in the solution? (Mole fraction of solute = 0.090; molar mass of NaOH = 40.0 g/mol; molar mass of H2O = 18.0 g/mol) a. 113.8 g b. 56.25 g c. 256.5 g d. 6.32 g e. 125.6 g a; Moderate 644. A solution is made by mixing 6.000 g of C10H8 in 90.00 g of C6H6. If the vapor pressure of pure C6H6 is 39.30 torr, what is the vapor pressure of the solution? (Molar masses: C10H8 = 128.0 g/mol; C6H6 = 78.00 g/mol)

a. 35.64 torr b. 1.201 torr c. 1.843 torr d. 37.76 torr e. 0.9609 torr d; Easy 645. The mole fraction of the solute in a solution is 0.1590. If the vapor pressure of the solvent in its pure state is 56.30 torr, what is the vapor pressure of the solution? a. 56.45 torr b. 47.35 torr c. 17.90 torr d. 49.95 torr e. 39.55 torr b; Easy 646. What is the boiling point of a 5.60 m solution of C6H4Cl2 in CCl4? Assume that C6H4Cl2 is not volatile. (Kb = 4.95°C/m for CCl4; TBP = 76.8°C) a. 105°C b. 101°C c. 89.8°C d. 91.2°C e. 123°C a; Easy 647. The boiling point of a solution of NaOH in H2O is 100.84°C. What is the molality of the solution? Assume that NaOH is not volatile. (Kb = 0.51200°C/m for H2O4; TBP = 100.00°C). a. 1.112 m b. 1.857 m c. 1.641m d. 0.9568 m e. 0.8205 m e; Easy 648. What is the freezing point of a 1.8 m solution of CBr4 in C6H6? (Kf for C6H6 = 4.90°C/m; Freezing point of C6H6 = 5.51°C) a. –6.90 °C b. 2.30 °C c. –3.31 °C d. 0.212 °C e. 1.20 °C c; Easy 649. What is the osmotic pressure of a 2.726 M solution of C6H12O6 at 47.00°C? a. 98.63 atm b. 71.57 atm c. 840.32 atm d. 10.13 atm e. 4.656 atm b; Easy

650. What is the ideal van‘t Hoff factor for LiNO3? a. 2 b. 3 c. 4 d. 5 e. 6 a; Easy 651. What is the van‘t Hoff factor for Fe(NO3)3? a. 11 b. 5 c. 2 d. 3 e. 4 e; Easy 652. What is the freezing point of a 10.0 m solution of NaOH in water? (Kf for water = 1.86 °C/m)? a. –18.6°C b. –3.20°C c. –11.6°C d. –37.2°C e. –9.25°C d; Easy 653. The osmotic pressure of a 0.0750 M KCl solution at 25.0°C is 3.28 atm. What is the true van‘t Hoff factor of this ionic compound? a. 3.28 b. 0.0750 c. 0.250 d. 3.48 e. 1.79 e; Easy 654. Determine the freezing point of a 2.33 m solution of PCl5 in H2O. (Kf = 1.86°C/m) a. –8.66°C b. 8.66°C c. –2.10°C d. –4.33°C e. 1.20°C d; Moderate 655. The osmotic pressure of blood is 4.83 atm at 47.0°C. If blood were considered a solution of NaCl, what is the molal concentration of NaCl in blood? Assume an ideal van‘t Hoff factor. a. 0.345 m b. 0.690 m c. 0.920 m d. 0.460 m e. 0.230 m

c; Easy Essay 656. How do you determine the solute and solvent in a solution? Explain with an example.

The major component of a solution is called the solvent. The minor component of a solution is called the solute. By major and minor we mean whichever component has the greater presence by mass or by moles.

Consider a solution is formed by 1 g of glucose and 100 g of water. Either by mass or by moles, the obvious minor component is glucose, so it is the solute. Water, the majority component, is the solvent.

Easy 657. Explain the concept of concentration in solutions.

The amount of a solute dissolved in a given amount of solvent is called concentration.

Dilute describes a solution that has very little solute, while concentrated describes a solution that has a lot of solute.

Easy 658. What is molarity? Explain with an example.

Molarity (M) is defined as the number of moles of solute divided by the number of liters of solution. moles of solution molarity = liters of solution A liter solution that contains 0.5 mol of solute has a molarity of 0.5 M. Easy 659. Explain the concept of molality.

Molality is similar to molarity. Molality is defined as the number of moles of solute per kilogram of solvent. moles of solution molarity = kilograms of solvent Easy 660. Explain the concentration units parts per thousand (ppth), parts per million (ppm), and parts per billion (ppb).

The concentration units are defined as follows. mass of solute ppth=  1000 mass of sample ppm=

ppb=

mass of solute  1,000,000 mass of sample

mass of solute  1,000,000,000 mass of sample

Easy 661. What is dilution and concentration?

Dilution is the addition of solvent, which decreases the concentration of the solute in the solution. Concentration is the removal of solvent, which increases the concentration of the solute in the solution. In both dilution and concentration, the amount of solute stays the same.

Easy 662. You are given 15 liters of a solution of NaCl in water. You estimate that the given solution has a molarity of 0.50 M. How do you increase the molarity of this solution to 1.0 M? The solution can be concentrated to 1 M by removing the solvent (water). moles of solute molarity (M) = liters of solution (V) Moles of the solute = MV In both dilution and concentration, the amount of solute stays the same. So MV is a constant.

Therefore M1V1 = M2V2 Here, 0.5 M × 15 L = 1 M × V2 V2 =

0.5 M ×15 L 1M

V2 = 7.5 L In order to concentrate the given solution to a 1 M solution, 7.5 L water should be removed from the solution.

Moderate 663. How many liters of a 0.5 M solution of CaCl2 will contain 10 moles of CaCl2? Explain your calculations.

moles of solute liters of solution (V) 10 mol 0.5= liters of solution 10 liters of solution =  20 L 0.5 molarity (M) =

Easy 664. What mass of solute is present in 1 L of 1.60 M solution of NaCl? Explain the calculations.

This is a two-step conversion, first using concentration as a conversion factor to determine the number of moles and then the molar mass of NaCl (58.45g/mol) to convert to mass:

1L 

1.60 mol NaCl 58.45g of NaCl = 93.5 g NaCl  L solution 1 mol of NaCl

Easy 665. What are colligative properties?

The properties of solutions are very similar to the properties of their respective pure solvents. This is because the majority of the solution is the solvent. However, some of the properties of solutions differ from pure solvents in measurable and predictable ways. The

differences are proportional to the fraction that the solute particles occupy in the solution.

These properties are called colligative properties.

Easy 666. Explain the concept of mole fraction with an example.

The mole fraction of the ith component in a solution, χi, is the number of moles of that component divided by the total number of moles in the sample. χi =

mol of i th component total moles

Easy 667. Explain Raoult‘s law.

The actual vapor pressure of a solution is calculated as follows: Psoln = χsolvP*solv.

where Psoln is the vapor pressure of the solution, χsolv is the mole fraction of the solvent particles, and P*solv is the vapor pressure of the pure solvent at that temperature (which is data that must be provided). This equation is known as Raoult‘s law.

Easy

668. Explain the concept of boiling point elevation.

Because the vapor pressure of a solution with a nonvolatile solute is depressed compared to that of the pure solvent, it requires a higher temperature for the solution‘s vapor pressure to reach 1.00 atm. The temperature at which the vapor pressure of the liquid equals 1.00 atm is called its boiling point. As such, the normal boiling point of the solution is higher than that of the pure solvent. This property is called boiling point elevation.

Easy 669. Explain freezing point depression.

The freezing point of a solution is lower than the freezing point of the pure solvent. Solute particles interfere with solvent particles coming together to make a solid, so it takes a lower temperature to get the solvent particles to solidify. This is called freezing point depression.

Easy 670. What is the reason why you should not drink seawater if you‘re stranded in a lifeboat on an ocean?

Sea water has a higher osmotic pressure than most fluids in your body. By ingesting seawater, it will pull the water out of your cells as osmosis works to dilute the seawater.

Your cells will die of thirst.

Moderate Fill in the Blanks 671. The minor component of a solution is called the _____. solute; Easy 672. A solution that has a lot of solute is called a(n) _____ solution.

concentrated; Easy 673. An unstable solution with more than normal maximum amount of solute in it is called a _______________ solution. supersaturated; Easy 674.

moles of solute  _____ . liters of solution molarity; Easy

675. The molarity is 3.50 M when _____ mol of NaCl is found in 0.500 L of solution.

1.75; Medium 676. _____ is defined as the number of moles of solute per kilogram of solvent.

Molality; Easy 677. _____ is the addition of solvent, which decreases the concentration of the solute in the solution.

Dilution; Easy 678. A 10.0 L of a 0.055 M solution are diluted to 15.8 L. The final concentration is _____ M. 0.035; Easy 679. _____ liters of 0.005 M NaOH are needed to obtain 1.00 mol of NaOH. 200; Easy 680. _____ g solute is present in 15 L of 0.20 M NaOH. (Molar mass of NaOH = 40.0 g/mol) 120; Easy 681. _____ is a property of solutions related to the fraction that the solute particles occupy in the solution, not their identity.

Colligative property; Easy

moles of i th component 682. _____ = total moles Mole fraction of the ith component in a solution; Easy 683. _____ is described by Psoln = χsolvP*solv.

Raoult’s law; Easy 684. _____ is the boiling point of a 6.200 m solution of NaCl in H2O. Assume that NaCl is not volatile (Kb for water = 0.5120 °C/m).

106.35 degrees Celsius; Easy 685. The ideal van‘t Hoff Factor for FeCl3 is _____. four; Easy Chapter 12 Acids and Bases

True/False 1. H+ ions and H3O+ ions are often considered interchangeable when writing chemical equations.

True; Easy 2. HCl is an Arrhenius acid.

True; Easy 3. NaCl is an Arrhenius base.

False; Easy 4. An Arrhenius acid is a compound that increases the hydroxide ion concentration in aqueous solution.

False; Easy 5. An Arrhenius acid is an ionic compound.

True; Easy 6. A reaction between an acid and a base is called a combustion reaction.

False; Easy 7. The Brønsted-Lowry definition of acid and base is limited to aqueous (that is, water) solutions.

False; Easy 8. When ammonia reacts with water, ammonia acts like a Brønsted-Lowry base.

True; Easy 9. In the reaction H3PO4 + OH−  H2PO4− + H2O, H3PO4 and H2PO4− are a conjugate acidbase pair.

True; Moderate 10. The titrant is added to the analyte using a precisely calibrated volumetric delivery tube called a buret.

True; Moderate 11. A base is always the titrant.

False; Moderate 12. HClO4 is a weak acid.

False; Easy 13. HCOOH is a strong acid.

False; Easy 14. LiOH is strong base.

True; Easy 15. Mg(OH)2 is a weak base.

False; Easy

16. Salts with ions that are part of weak acids or bases will not hydrolyze, while salts with ions that are part of strong acids or bases will hydrolyze.

False; Moderate 17. The product of the two concentrations—[H+][OH−]—is always equal to 1.0 × 10−14.

True; Easy 18. Consider the product of the two concentrations, [H+][OH−]. If one concentration goes up, the other must go up to compensate so that their product equals the value of the autoionization constant of water.

False; Moderate 19. If pOH < 7, then the solution is acidic.

False; Moderate 20. If pH > 7, then the solution is basic.

True; Moderate 21. Pure water (pOH = 7) is an acidic solution.

False; Moderate 22. pH + pOH = 14 True; Easy 23. [OH−] = 10−pOH + pH False; Moderate 24. Small amounts of weak acids and bases can change the pH of a solution very quickly.

False; Easy 25. HN3 and NH3 can make a buffer.

False; Moderate Multiple Choice Questions 26. Chemistry defines the _____ ion as the actual chemical species that represents an H+ ion. z. phosphate aa. sulfate bb. nitrate cc. nitrite dd. hydronium e; Easy 27. Which of the following is an Arrhenius acid? a. H2SO3 b. NaOH c. Na2CO3 d. HgI2 e. Al2(SO4)3 a; Easy

28. Which of the following is not an Arrhenius acid? a. H2SO3 b. H2SO4 c. HClO4 d. HI e. Li2SO4 e; Easy 29. Which of the following is an Arrhenius base? a. HC2H3O2 b. HClO3 c. Al(OH)3 d. PCl5 e. HgCl2 c; Easy 30. Which of the following statements is true about acids? a. Acids turn litmus, a plant extract, red. b. Acids react with hydrogen carbonate to give of N2 gas. c. Acids that are ingested typically have a sweet taste. d. Acids are slippery to touch. e. Acids increase the concentration of hydroxide ions in a solution. a; Moderate 31. A reaction between an acid and a base is called a(n): a. combustion reaction. b. decomposition reaction. c. hydrogenation reaction. d. neutralization reaction. e. alkylation reaction. d; Easy 32. Which of the following is formed when H2SO3 reacts with KOH? a. K2SO4 b. K2SO3 c. K2(SO4)3 d. KH e. KOH b; Moderate 33. Which of the following is required to react with Zn(OH)2 to produce Zn3(PO4)2? a. H2PO3 b. H2PO4 c. HPO4 d. H3PO4 e. H2P2O4 d; Moderate 34. Complete the following reaction. ? + 2HNO3  Sr(NO3)2 + 2H2O a. Sr(OH)2

b. SrNO3 c. Sr2NO3 d. SrOH2 e. SrON3 a; Moderate 35. A Brønsted-Lowry acid is any species that can donate a(n) _____to another molecule. a. hydroxide (OH–) ion b. water (H2O) molecule c. proton (H+) d. ammonium (NH4+) ion e. electron (e-) c; Easy 36. Complete the following reaction. ? + H2O  C6H5COO– + H3O+ a. C6H5COOH b. C7H8COOH c. C6H13COOH d. C5H11COOH e. C7H15COOH a; Moderate 37. Predict the products of this reaction, assuming it undergoes a Brønsted-Lowry acid-base reaction.

H2C2O4 + 2F−  ? a. 2H2C2O4 and F− b. C2O42− and 2HF c. C2O4 and 2F− d. H2C2O3– and F2 e. C2O3– and 2F2 b; Moderate 38. Predict the products of this reaction, assuming it undergoes a Brønsted-Lowry acid-base reaction.

HCOOH + C5H5N  ? a. HCOON and C6H6 b. HCOO– and C5H5NH+ c. HCOOH2+ and C5H4N– d. H2COOH2+ and C5H3N– e. CHCOOH and C4H5N– b; Moderate 39. What is the conjugate acid of H2O? a. H2O+ b. H2O2– c. HO3– d. HO3+ e. H3O+ e; Easy

40. What is the conjugate acid of ClO4–? a. HClO32– b. ClO4– c. HClO4 d. H2ClO4+ e. H2ClO4– c; Easy 41. What is the conjugate base of C5H5NH+? a.

C5H5NH2+ b.

C5H5NH– c.

C5H5NH2+ d.

C5H5NH2 e. C5H5N e; Easy 42. What is the conjugate base of HClO3? a.

HClO3– b.

HClO3+ c.

ClO3– d.

H2ClO3– e.

ClO4– c; Easy

43. What mass of AgOH is present in a sample if it is titrated to its equivalence point with 58.5 mL of 0.0663 M HNO3? The balanced chemical equation is as follows: AgOH + HNO3  AgNO3 + H2O a. 0.485 g b. 0.00388 g c. 258 g d. 1.13 g e. 2.27 g a; Moderate 44. What mass of Ca(OH)2 is present in a sample if it is titrated to its equivalence point with 18.09 mL of 0.2235 M H2C2O4? The balanced chemical reaction is as follows: H2C2O4 + Ca(OH)2 → CaC2O4 + 2H2O a. 248.0 g b. 0.2992 g c. 0.004043 g d. 0.2235 g e. 0.5984 g b; Moderate 45. 32.66 mL of 0.6664 M Zn(OH)2 was needed to titrate a sample of 0.6664 M HC2H3O2 to its equivalence point. Determine the initial volume of the HC2H3O2 solution. a. 16.30 mL b. 32.70 mL c. 65.30 mL d. 24.50 mL e. 10.00 mL

c; Moderate 46. If 100.0 mL of 0.2221 M HCl is needed to titrate a sample of NaOH to its equivalence point, the mass of NaOH present in the sample is _____. a. 0.6663 g b. 0.2221 g c. 0.4442 g d. 1.717 g e. 0.8884 g e; Moderate 47. 30.33 mL of 0.7774 M HC2H3O2 is needed to titrate a sample of 28.45 mL KOH to its equivalence point. What is the original concentration of the KOH solution? a. 0.4144 M b. 0.2072 M c. 2.486 M d. 0.8288 M e. 1.658 M d; Moderate 48. Which of the following acids ionizes 100% in aqueous solution? a. HBr b. HBrO3 c. HNO2 d. HCOOH e. CH3COOH a; Easy 49. Which of the following is a weak acid? a. HCl b. H2SO4 c. H3PO4 d. HI e. HClO4 c; Easy 50. Which of the following is a strong base? a. Sr(OH)2 b. Hg(OH)2 c. Mn(OH)2 d. Fe(OH)2 e. Fe(OH)3 a; Easy 51. Which of the following bases does not ionize 100% in aqueous solution? a. LiOH b. Mg(OH)2 c. KOH d. Ca(OH)2 e. Ni(OH)2 e; Easy

52. Which of the following is a basic salt? a. LiCl b. RbBr c. KNO2 d. NaI e. NH4Cl c; Moderate 53. Which of the following is an acidic salt? a. RbCl b. Sr(NO2)2 c. CaNO2 d. CaBr2 e. MnI2 e; Moderate 54. Which of the following is a neutral salt? a. Al2(SO4)3 b. Ti(NO3)2 c. Ti(PO4)3 d. Fe(ClO3)2 e. Ba(NO3)2 e; Moderate 55. The concentration of both H+(aq) and OH−(aq) in a sample of pure H2O is about: a. 1.0  107 M. b. 1.0  10–7 M. c. 1.0  10–14 M. d. 1.0  10–10 M. e. 1.0  10–21 M. b; Easy 56. What is [OH−] of an aqueous solution if [H+] is 1.0 × 10−7 M? a. 1.0  10–14 M b. 1.0  10–5 M c. 1.0  105 M d. 1.0  10–7 M e. 1.0  10–21 M d; Easy 57. What is [H+] of an aqueous solution if [OH−] is 1.0 × 10−5 M? a. 1.0  10–8 M b. 1.0  10–14 M c. 1.0  107 M d. 1.0  10–7 M e. 1.0  10–9 M e; Easy 58. What is [H+] in a 0.0066 M solution of Ba(OH)2?

a. 7.6  10–13 M b. 1.0  1014 M c. 1.0  10–13 M d. 1.5  10–12 M e. 0.0066 M a; Moderate 59. What is [OH–] in a 0.077 M solution of HClO4? a. 7.6  10–13 M b. 1.3  10–13 M c. 6.5  10–14 M d. 6.5  1014 M e. 0.077 M b; Moderate 60. A 0.0788 M solution of HC2H3O2 is 3.0% ionized into H+ ions and C2H3O2− ions. What is [OH−] for this solution? a. 8.5  10–12 M b. 2.1  10–12 M c. 4.2  10–12 M d. 8.5  1012 M e. 0.0788 M c; Moderate 61. A 0.0111 M solution of Sn(OH)2 is 0.585% ionized into Sn2+ and OH– ions. What is [H+] for this solution? a. 1.54  10–10 M b. 7.70  10–11 M c. 3.08  10–10 M d. 7.70  1011 M e. 0.0111 M b; Moderate 62. What is [OH−] for an aqueous solution whose pH is 3.89? a. 7.76  10–11 M b. 1.29  10–12 M c. 1.29  1012 M d. 2.58  1012 M e. 3.86  10–12 M a; Moderate 63. What is [H+] for an aqueous solution whose pOH is 9.89? a. 2.33  10–6 M b. 7.76  10–5 M c. 2.33  106 M d. 4.11 M e. 1.55  10–6 M b; Moderate

64. What is the pH of a solution when [OH–] is 8.04 × 10−13 M? a. 0.904 b. 1.91 c. 0.975 d. 14.0 e. 12.09 b; Moderate 65. What is the pOH of a solution when [H+] is 5.88 × 10−11 M? a. 10.1 b. 7.88 c. 10.23 d. 3.77 e. 15.6 d; Moderate 66. Which of the following combinations of compounds can make a buffer solution? a. HCOOH and CH3COOH b. NaCl and BaCl2 c. HCl and Ca(OH)2 d. H3PO4 and Li3PO4 e. NH3 and NaOH d; Moderate 67. Which of the combinations of compounds can make a buffer solution? a. Ca(HCO3)2 and Ca(OH)2 b. CH3NH2 and CH3NH3Br c. NaOH and NaCl d. HCl and NaCl e. NH3 and Ca(OH)2 b; Moderate 68. Which of the combinations of compounds can make a buffer solution? a. HCl and LiCl b. HNO2 and HF c. RbOH and CsOH Rb2SO4 and HNO3 d. e. NH4NO3 and NH3 e; Moderate 69. Which of the combinations of compounds can make a buffer solution? a. LiOH and LiCl b. NaOH and NaBr NaNO3 and Ca(NO3)2 c.

NaHCO3 and Na2CO3 d.

HN3 and NH3 e. d; Moderate

70. Which of the following can be a component in a buffer? a. LiBr b. NaI

c. KNO3 d. K3PO4 e. SrI2 d; Moderate Essay Questions 71. Differentiate between Arrhenius acids and Arrhenius bases.

An Arrhenius acid is a compound that increases the H+ ion concentration in aqueous solution.

An Arrhenius base is a compound that increases the OH– ion concentration in aqueous solution.

Easy 72. Explain neutralization reactions with examples.

A reaction between an acid and a base is called a neutralization reaction and can be represented as follows: acid + base → H2O + salt The stoichiometry of the balanced chemical equation depends on the number of H+ ions in the acid and the number of OH− ions in the base.

E.g.: The neutralization reaction between H2SO4 and KOH is as follows: Because the acid has two H+ ions in its formula, we need two OH− ions to react with it, making two H2O molecules as product. The remaining ions, K+ and SO42−, make the salt potassium sulfate (K2SO4). The balanced chemical reaction is as follows: H2SO4 + 2KOH → 2H2O + K2SO4 (Students examples may vary.)

Moderate 73. Differentiate Brønsted-Lowry acids and Brønsted-Lowry bases.

A Brønsted-Lowry acid is any species that can donate a proton (H+) to another molecule.

A Brønsted-Lowry base is any species that can accept a proton from another molecule.

In short, a Brønsted-Lowry acid is a proton donor (PD), while a Brønsted-Lowry base is a proton acceptor (PA).

Easy 74. Why is ammonia unique from the other bases? Describe the hydrolysis of ammonia.

Ammonia (NH3) is a base even though it does not contain OH− ions in its formula.

Instead, it generates OH− ions as the product of a proton-transfer reaction with H2O molecules; NH3 acts like a Brønsted-Lowry base, and H2O acts like a Brønsted-Lowry acid:

A reaction with water is called hydrolysis; we say that NH3 hydrolyzes to make NH4+ ions and OH− ions.

Moderate 75. Briefly describe conjugate acid-base pairs.

In the reaction between NH3 and H2O, two sets of similar species exist on both sides.

Within each set, the two species differ by a proton in their formulas, and one member of the set is a Brønsted-Lowry acid, while the other member is a Brønsted-Lowry base.

The two sets—NH3/NH4+ and H2O/OH−—are called conjugate acid-base pairs.

Moderate 76. Describe amphiprotic substances with an example.

A substance that can act as a proton donor or a proton acceptor is called amphiprotic.

Water is the most common amphiprotic substance.

When water reacts with ammonia, water acts as an acid and donates a proton to ammonia.

NH3(g) + H2O(l)  NH4+(aq) + OH–(aq) When HCl(g) is dissolved in water, HCl(g) acts as the proton donor, while H2O is the proton acceptor.

HCl(g) + H2O(ℓ) → H3O+(aq) + Cl−(aq) Moderate 77. Describe the process of titration.

Performing chemical reactions quantitatively to determine the exact amount of a reagent is called a titration.

In a titration, one reagent has a known concentration or amount, while the other reagent has an unknown concentration or amount. Typically, the known reagent (the titrant) is added to the unknown quantity and is dissolved in solution. The unknown amount of substance (the analyte) may or may not be dissolved in solution (but usually is). The titrant is added to the analyte using a precisely calibrated volumetric delivery tube called a burette. The burette has markings to determine how much volume of solution has been added to the analyte. When the reaction is complete, it is said to be at the equivalence point; the number of moles of titrant can be calculated from the concentration and the volume, and the balanced chemical equation can be used to determine the number of moles (and then concentration or mass) of the unknown reactant.

Moderate 78. Differentiate between strong acids and weak acids.

Any acid that dissociates 100% into ions is called a strong acid. If it does not dissociate 100%, it is a weak acid.

Easy 79. Differentiate between strong bases and weak bases. Give an example of two strong acids and two strong bases.

Any base that dissociates 100% into ions is called a strong base. If it does not dissociate 100%, it is a weak base. Examples can include (strong acids) HCl, HBr, HI, H2SO4, HNO3, HClO3, HClO4. Strong bases: LiOH, NaOH, KOH, RbOH, CsOH, Mg(OH)2, Ca(OH)2, Sr(OH)2, Ba(OH)2.

Moderate 80. What are the rules required to determine whether a solution is acidic or basic?

There are two general rules: (1) If an ion derives from a strong acid or base, it will not affect the acidity of the solution.

(2) If an ion derives from a weak acid, it will make the solution basic; if an ion derives from a weak base, it will make the solution acidic.

Moderate 81. Briefly describe autoionization of water. What is the autoionization constant of water?

Some H2O molecules act as acids, and other H2O molecules act as bases. The chemical equation is as follows: H2O + H2O → H3O+ + OH− This is called autoionization of water.

In acids, the concentration of H+(aq)—[H+]—is greater than 1.0 × 10−7 M, while for bases the concentration of OH−(aq)—[OH−]—is greater than 1.0 × 10−7 M. However, the product of the two concentrations—[H+][OH−]—is always equal to 1.0 × 10−14.

[H+][OH−] = 1.0 × 10−14 This value of the product of concentrations is so important for aqueous solutions that it is called the autoionization constant of water.

Moderate 82. Explain how the pH scale can be used to determine whether a given aqueous solution is acidic, basic, or neutral. pH is a logarithmic function of [H+]: pH = −log[H+] If pH < 7, then the solution is acidic.

If pH = 7, then the solution is neutral.

If pH > 7, then the solution is basic.

This is known as the pH scale. You can use pH to make a quick determination whether a given aqueous solution is acidic, basic, or neutral.

Moderate 83. Demonstrate how a buffer works.

Let us use an HC2H3O2/NaC2H3O2 buffer to demonstrate how buffers work. If a strong base—a source of OH−(aq) ions—is added to the buffer solution, those OH− ions will react with the HC2H3O2 in an acid-base reaction: HC2H3O2(aq) + OH−(aq) → H2O(ℓ) + C2H3O2−(aq)

Rather than changing the pH dramatically by making the solution basic, the added OH− ions react to make H2O, so the pH does not change much.

If a strong acid—a source of H+ ions—is added to the buffer solution, the H+ ions will react with the anion from the salt. Because HC2H3O2 is a weak acid, it is not ionized much. This means that if lots of H+ ions and C2H3O2− ions are present in the same solution, they will come together to make HC2H3O2: H+(aq) + C2H3O2−(aq) → HC2H3O2(aq) Rather than changing the pH dramatically and making the solution acidic, the added H+ ions react to make molecules of a weak acid.

Moderate 84. Explain why KNO3 cannot be a component in either an acidic or a basic buffer.

KNO3 is a salt that is derived from a strong acid and a strong base, so it is not a salt of either a weak acid or weak base and cannot be a component of a buffer.

Easy 85. Two solutions are made containing the same concentrations of solutes. One solution is composed of NaHCO3 and Na2CO3, while the other is composed of HCN and NaCN. Which solution should have the larger capacity as a buffer?

Neither has the larger capacity. They can accept or donate only a single H+ or OH−.

Easy Fill in the Blanks 86. A(n) _____ is a compound that increases the hydrogen ion concentration in aqueous solution.

Arrhenius acid; Easy 87. A(n) _____ is a compound that increases the hydroxide ion concentration in aqueous solution.

Arrhenius base; Easy 88. A _____is any ionic compound made by combining an acid with a base. salt; Easy 89. A reaction between an acid and a base is called a(n)_____ reaction. neutralization; Easy 90. A _____is any species that can accept a proton (H+) from another molecule.

Brønsted-Lowry base; Easy 91. When ammonia reacts with water, water acts like a_____.

Brønsted-Lowry base; Easy 92. A substance that can act as a proton donor or a proton acceptor is called _____. amphiprotic; Moderate 93. Performing chemical reactions quantitatively to determine the exact amount of a reagent is called a(n)_____. titration; Easy

94. The titrant is added to the analyte using a precisely calibrated volumetric delivery tube called a(n)_____. burette; Easy 95. A(n) _____ is a substance that changes color depending on the acidity or basicity of the solution. indicator; Easy 96. RbCl is an example of a(n)_____ salt. neutral salt; Moderate 97. Water molecules act as acids (proton donors) and bases (proton acceptors) with each other to a tiny extent in all aqueous solutions. This is known as the _____. autoionization of water; Moderate 98. The product of the two concentrations—[H+][OH−]—is always equal to 1.0 × 10−14. This is known as the _____. autoionization constant of water; Moderate 99. If pH is _____, then the solution is acidic. < 7; Easy 100. A(n) _____ is a solution that resists dramatic changes in pH. buffer; Easy 101. Buffers work well only for limited amounts of added strong acid or base. Once either solute is completely reacted, the solution is no longer a buffer, and rapid changes in pH may occur. We say that a buffer has a(n) ______. capacity; Easy Chapter 13 Chemical Equilibrium True/False 1. An equilibrium equation should be balanced.

True; Easy 2. The equilibrium equation that exists between calcium sulfate and water as reactants and calcium hydroxide and sulfuric acid as products can be represented as True; Easy 3. Equilibrium reactions do not stop; both the forward and reverse reaction continue to occur.

True; Moderate 4. Chemical equilibrium is static rather than dynamic.

False; Moderate 5. For a given reaction

, the equilibrium constant can be defined as

a

b

é A ù é Bù K eq = ë ûc ë ûd . éëC ùû éë D ùû False; Easy 6. In a chemical equilibrium reaction all the three phases of matter are included. This equation is in a state of homogeneous equilibrium.

False; Moderate 7. Given: If [HCl], [Cl2], and Keq are 0.586 M, 0.486 M, and 2.37, respectively, the equilibrium [H2] is 3.35 M. False; Moderate 8. Given: If [C2H6], [O2], [CO2], and Keq are 0.451 M, 0.589 M, 0.638 M, and 25.0, respectively, the equilibrium [H2O] is 0.632 M. False; Moderate 9. Given: If [C2H6], [O2], [CO2], and Keq are 0.451 M, 0.589 M, 0.638 M, and 25.0, respectively, the equilibrium [H2O] is 0.326 M. True; Moderate 10. Given: The equilibrium partial pressures of [H2O2] and [O2] are 1.45 atm and 1.89 atm, respectively. If KP = 6.00 the equilibrium partial pressure of H2O is 0.134 atm.

False; Moderate 11. Given: If the Keq at 57.0°C for the reaction above is 13.6, then KP is 0.0474.

Assume R = 0.08205

L  atm mol  K

False; Moderate

12. Given: To maximize the amount of N2O4 in the reaction above, the number of moles of NO2 should be increased.

True; Moderate 13. Given: To maximize the reactants in the reaction above, the temperature of the reaction should be decreased.

False; Moderate

14. Given: Decreasing the pressure in the reaction above shifts the equilibrium toward the reactants.

False; Moderate 15. Given: Increasing the temperature in the reaction above shifts the equilibrium toward the reactants.

True; Moderate 16. A catalyst does not affect the extent or position of a reaction at equilibrium.

True; Easy 17. The equilibrium concentration of H2S(g) in the reaction below is 0.17 M.

False; Moderate 18. The equilibrium concentration of H2O(g) for the chemical reaction below is 0.900 M.

False; Moderate 19. The Ka for HClO2 is 1.10 × 10−2. The pH of 0.165 M HClO2 in H2O is 1.37.

True; Moderate 20. The Ka for HCN is 6.20 × 10−10. The concentration of CN– in the ionization of 0.0400 M HCN in H2O is 4.89 × 10−6 M False; Moderate 21. If the Kb for CN− is 1.6 × 10−5, Ka for HCN is 6.3 × 10−10.

True; Moderate 3

SO 24  22. The expression for the slight solubility of Al2(SO4)3(s) is K sp  .

2  Al3  False; Moderate 23. If [OH−] is 2.02 × 10−7 M, [H+] in a solution is 4.59 × 10−8.

False; Moderate 24. If the Ksp of SrSO4(s) is 3.8 × 10−4, the concentration of [SO42–] in a saturated solution of SrSO4(s) is 9.1 × 10−12 M. True; Moderate

Multiple Choice Questions 25. Write the equilibrium equation that exists between hydrogen and chlorine as reactants and hydrogen chloride as a product. u. v. w. x. y. e; Easy 26. Write the equilibrium equation that exists between calcium sulfate as a reactant and calcium oxide and sulfur trioxide as products. a. b. c. d. e. c; Easy Refer to the reaction below for Multiple Choice Questions 28-29.

Given: 27. What is the Keq expression for the reaction? a.

K eq =

éë HBr ùû éë H 2 ùû éë Br2 ùû 2

é HBr ùû b. K eq = ë 2 éë H 2 ùû éë Br2 ùû

 HBr  K eq  2  H 2   Br2  2 HBr   K eq   H 2  Br2   HBr  K eq  2  H2  Br2  2

c.

d. e.

d; Easy 28. If [H2], [Br2], and [HBr] are 0.0350 M, 0.0600 M, and 0.0850 M, respectively, determine Keq.

a. 6.00 b. 4.50 c. 3.44 d. 1.80 e. 404 c; Moderate Refer to the reaction below for Multiple Choice Questions 30-31.

Given: 29. What is the Keq expression for the reaction? a.

b.

c.

 HCl  H 2 Cl2  2 HCl  K eq  2  H 2 Cl2   H Cl2  K  2 K eq 

eq

[HCl]2

 H Cl  d. K  2

eq

e.

K eq 

2

2

[HCl]  H2 Cl2 

[HCl]

c; Easy 30. If [H2], [Cl2], and Keq are 0.315 M, 0.218 M, and 3.85, respectively, determine the equilibrium [HCl]. a. 0.240 M b. 0.514 M c. 0.0354 M d. 0.134 M e. 0.0530 M d; Moderate Refer to the reaction below for Multiple Choice Questions 32-33.

Given: 31. What is the Keq expression for the reaction?

CO2   H 2O K eq  4 2  O2  C2 H 2  4 2 CO 2   H 2O   K eq  2 5  O2  C2 H 2  5

a.

b.

2

CO2   H 2O K eq  5 2  O2   C2 H 2  4 2 CO 2   H 2O   K eq  5 2  O2   C2 H 2  4 2 O2  C2 H 2   K eq  5 2 CO2   H 2O 2

c.

d.

e.

4

d; Moderate 32. If [C2H2], [O2], [H2O], and [CO2] are 0.295 M, 0.400 M, 0.500 M, and 0.686 M, respectively, determine Keq. a. 84.4 b. 48.4 c. 62.3 d. 42.2 e. 22.4 c; Moderate Refer to the reaction below for Multiple Choice Questions 34-35.

Given: 33. What is the Keq expression for the reaction?

CO2   H 2O K eq  4 2  O2  C2 H 2  4 2 CO 2   H 2O   K eq  2 5  O2  C2 H 2  2 4 CO 2   H 2O   K eq  5 2  O2   C2 H 2  4 2 CO 2   H 2O   K eq  5 2  O2   C2 H 6  2 7 C2 H 6   O2   K eq  6 4  H 2O CO2  5

a.

b.

c.

d.

e.

2

e; Moderate 34. If [C2H6], [O2], [CO2], and Keq are 0.391 M, 0.486 M, 0.400 M, and 25.0, respectively, determine the equilibrium [H2O]. a. 0.993 M b. 0.339 M c. 0.393 M d. 0.933 M

e. 0.399 M b; Moderate Refer to the reaction below for Multiple Choice Questions 36-37.

Given: 35. What is the KP expression for the reaction? a.

KP 

b.

KP 

c.

KP 

d.

KP 

e.

KP 

PCl2 PPCl3  PPCl5 PPCl3 PPCl5  PCl2 PPCl5 PCl2  PPCl3 PCl2  PPCl3 PPCl5

PCl2  PPCl3 PPCl5

c; Easy 36. If the equilibrium partial pressures of [PCl5], [PCl3], and [Cl2] are 0.240 atm, 0.291 atm, and 0.655 atm, respectively, determine KP. a. 3.14 b. 1.43 c. 1.26 d. 4.31 e. 4.13 c; Easy Refer to the reaction below for Multiple Choice Questions 38-39.

Given: 37. What is the expression for PO2 for this reaction? a.

PO2 

PH22O  K P PH22O2

 PH2 O  b. PO2   2 2 2   PH O  K P   2  2 PH O c. PO2  2 2 2 PH2O  K P

2

d.

e.

PO2 

PO2 

PH22O2 PH2O  K P2

PH22O2 PH2O  K P

c; Moderate 38. The equilibrium partial pressures of [H2O2] and [O2] are 0.239 atm and 0.281 atm, respectively. If KP = 5.00, determine the equilibrium partial pressure of [H2O]. a. 2.20 atm b. 0.220 atm c. 0.0022 atm d. 0.022 atm e. 0.202 atm e; Moderate 39. What is the KP at 37.0°C for this reaction if the Keq is 1.20 × 10−2?

Assume R = 0.08205 a. 4.27  10– 4 b. 2.47  10– 4 c. 2.74  10– 4 d. 7.42  10– 4 e. 4.72  10– 4 e; Moderate

L  atm mol  K

40. What is the KP at 47.0°C for this reaction if the Keq is 2.36?

Assume R = 0.08205

L  atm mol  K

a. 31.2 b. 9.88 c. 0.0988 d. 62.0 e. 0.0889 d; Moderate 41. What is the Keq at 7.00°C for this reaction if the KP is 0.00123?

Assume R = 0.08205 a. 19.4 b. 14.9 c. 49.1

L  atm mol  K

d. 1.94 e. 1.49 b; Moderate

42. What is the correct Keq expression for this reaction?

a. b.

c.

d.

e.

 HCl NaOH   NaCl H 2O  HCl NaOH  K eq  2  NaCl  H 2O  HCl NaOH  K eq   NaCl 2 HCl NaOH   K eq   NaCl 2 HCl  NaOH   K eq   NaCl K eq 

c; Moderate 43. What is the correct expression for the concentration of [C2H2I4] in this reaction?

 C2 H 2  2 K eq   I 2   C2 H 2  b.  C2 H 2 I 4   2 K eq  I2  c.  C2 H 2 I 4    C2 H 2   K eq C2 H 2  d.  C 2 H 2 I 4   a.

C2 H 2 I4  

K eq

e.

 C2 H 2 I4  

K eq

C2 H 2 

d; Moderate 44. What is the correct KP expression for this reaction?

a.

KP 

2 PCH 2 Cl 2

PCH 4  PCl2 2

b.

KP 

c.

KP 

d.

KP 

e.

KP 

PCH2Cl2 2 PCH  PCl2 2 4 2 PCH2Cl2  PHCl

PCH4  PCl2 2 2 PCH2Cl2  PHCl

PCH4  PCl2 2

PCH2Cl2 PCH4  PCl2 2

e; Moderate Refer to the reaction below for Multiple Choice Questions 46-49.

Given: 45. Which of the following will cause the equilibrium to shift toward the products? a. Addition of N2 b. Addition of H2 c. Removal of NH3 d. Increasing the pressure e. Increasing the temperature e; Moderate 46. Which of the following will cause the equilibrium to shift toward the reactants? a. Removal of N2 b. Removal of H2 c. Removal of NH3 d. Reducing the pressure e. Increasing the temperature c; Moderate 47. To maximize the amount of NH3 the: a. pressure of the reaction should be reduced. b. temperature of the reaction should be reduced. c. number of moles of H2 in the reaction should be reduced. d. number of moles of N2 in the reaction should be reduced. e. number of moles of NH3 in the reaction should be increased. b; Moderate 48. To maximize the amount of N2 the: a. pressure of the reaction should be reduced. b. temperature of the reaction should be decreased. c. amount of catalyst added should be increased. d. number of moles of N2 in the reaction should be increased. e. number of moles of NH3 in the reaction should be reduced. a; Moderate

Refer to the reaction below for Multiple Choice Questions 50-53.

Given: 49. Which of the following will cause the equilibrium to shift toward the reactants side? a. Addition of CO b. Removal of CO2 c. Removal of the catalyst d. Increasing the temperature e. Increasing the pressure d; Moderate 50. Which of the following will cause the equilibrium to shift toward the products side? a. Removal of CO b. Increasing the pressure c. Addition of CO2 d. Removal of the catalyst e. Increasing the temperature b; Moderate 51. To maximize the amount of CO the: a. pressure of the reaction should be reduced. b. temperature of the reaction should be reduced. c. number of moles of CO2 in the reaction should be reduced. d. amount of catalyst added should be increased. e. number of moles of CO in the reaction should be increased. a; Moderate 52. To maximize the amount of CO2 the: a. pressure of the reaction should be reduced. b. temperature of the reaction should be increased. c. number of moles of CO2 in the reaction should be increased. d. amount of catalyst added should be reduced. e. number of moles of CO in the reaction should be increased. e; Moderate 53. Which of the following is the right ICE chart for the reaction below?

a.

I C E

2O3(g) –2x –2x – 2x

3O2(g) 3x +3x +3x

2O3(g) 0.025 –x

3O2(g) 0 +x

b.

I C

E

0.025 – 2x

+3x

2O3(g) 0.025 –x 0.025 – x

3O2(g) 0 +x +x

2O3(g) 0.025 +x 0.025 + x

3O2(g) 0 –x –x

2O3(g) 0.025 –2x 0.025 – 2x

3O2(g) 0 +3x +3x

c.

I C E d.

I C E e.

I C E

e; Moderate 54. Determine the equilibrium concentration of H2(g) for this reaction.

a. 0.039 M b. 0.0039 M c. 0.39 M d. 0.093 M e. 0.0093 M b; Moderate 55. Determine the equilibrium concentration of CO(g) for this reaction.

a. 0.836 M b. 0.683 M c. 0.386 M d. 0.863 M e. 0.368 M d; Moderate 56. Determine the equilibrium concentration of Cl2(g) for this reaction.

a. 0.729 M b. 9.27 M c. 2.97 M d. 9.72 M e. 0.297 M c; Moderate 57. Determine the equilibrium concentration of N2O3(g) for this reaction.

a. 1.76 M b. 1.23 M c. 5.61 M d. 1.65 M e. 6.51 M b; Moderate 58. Determine the equilibrium concentration of HCN(g) for this reaction.

a. 3.27 M b. 2.37 M c. 0.372 M d. 0.370 M e. 0.730 M e; Moderate 59. If concentration is defined as x, formulate Keq for the reaction given below.

 0.075  3x  

a.

K eq

b.

K eq 

c.

K eq 

d.

K eq 

e.

K eq 

2

4 x3 4 x4

 0.075  3x 

3

4 x2

 0.075  3x 

3

4 x3

 0.075  3x 

2

3x 2

 0.075  4 x 

c; Moderate

3

60. If concentration is defined as x, formulate Keq for the reaction given below.

x2 x 2  0.049 x  0.00033 x2 b. K eq   2 x  0.061x  0.00033 x2 c. K eq  2 x  0.049 x  0.00033 x2 d. K eq   2 x  0.061x  0.00033 x2 e. K eq  2 x  0.061x  0.00033 a.

K eq 

e; Moderate 61. Which of the following is the equilibrium equation for HF acting as a weak acid? a. b. c. d. e.

d; Easy 62. Which of the following is the Ka expression for HSO4– acting as a weak acid?

2

a.

 H +   SO 24  Ka     2   HSO 4 

2

 H +   SO42  b. K a   HSO 4  2

c.

 H +   SO 24  Ka   HSO 4 

2

OH    SO32  d. K a   HSO 4  2

e.

 H +   SO32  Ka        HSO 4 

b; Moderate

2

2

63. What is the pOH of a 0.500 M solution of HClO2? The Ka of HClO2 is 1.10 × 10−2. a. 1.13 b. 12.9 c. 3.11 d. 11.3 e. 31.1 b; Moderate

64. What is the pH of a 0.600 M solution of H2PO4−? The Ka of H2PO4− is 6.2 × 10−8. a. 3.71 b. 10.3 c. 3.10 d. 30.1 e. 17.7 a; Moderate 65. What is the value of Kb for CN–, which can accept a proton and act as a base? The Ka for HCN is 6.20 × 10−10. a. 6.20 × 10−10 b. 2.60 × 10−5 c. 1.16 × 10−5 d. 6.11 × 10−5 e. 1.61 × 10−5 e; Easy 66. What is the Ksp expression for Mg3(PO4)2? a.

3

Ksp   Mg 2+   PO34 

2

3

 Mg 2+  b. K sp  2  PO34 

3

c.

 PO34  K sp  2  Mg 2+ 

d.

Ksp   PO34    Mg 2+ 

e.

Ksp   PO34    Mg 2+ 

3

2

3

2

a; Moderate 67. What is the concentration of [Ag+] in a saturated solution of AgI? The Ksp of AgI is 9.50 × 10−17. a. 9.75 × 10−9 M b. 2.92 × 10−10 M c. 6.20 × 10−10 M d. 2.60 × 10−9 M e. 2.29 × 10−9 M a; Easy

68. What is the concentration of [Ag+] in a saturated solution of Ag2SO4? The Ksp of Ag2SO4 is 1.50 × 10−5. a. 0.0155 M b. 0.0310 M c. 0.0465 M d. 0.0620 M e. 0.0775 M b; Moderate 69. What is the concentration of [PO43–] in a saturated solution of Ca3(PO4)2? The Ksp of Ca3(PO4)2 is 2.1 × 10−33. a. 3.28 × 10−6 M b. 1.64 × 10−6 M c. 2.82 × 10−7 M d. 2.28 × 10−7 M e. 8.22 × 10−7 M d; Moderate Essay Questions 70. Define chemical equilibrium with an example.

Consider the following reaction occurring in a closed container H2 + I2  2HI The way the equation is written, we are led to believe that the reaction goes to completion, that all the H2 and the I2 react to make HI. However, this is not the case. The reverse chemical reaction is also taking place: 2HI H2 + I2 The reverse reaction proceeds so quickly that it matches the speed of the forward reaction. When that happens, any continued overall reaction stops: the reaction has reached chemical equilibrium, the point at which the forward and reverse processes balance each other‘s progress.

Because two opposing processes are occurring at once, it is conventional to represent an equilibrium using a double arrow, like this: Easy 71. Describe chemical equilibrium as a dynamic process.

In equilibrium reactions the reactions do not stop; both the forward reaction and the reverse reaction continue to occur. They both occur at the same rate, so any overall change by one reaction is cancelled by the reverse reaction. We say that chemical equilibrium is dynamic, rather than static. Also, because both reactions are occurring simultaneously, the equilibrium can be written backward. For example, representing an equilibrium as Easy 72. Explain the law of mass action and equilibrium constant.

Law of mass action relates the amounts of reactants and products at equilibrium for a chemical reaction. For a general chemical reaction occurring in solution,

the equilibrium constant, also known as Keq, is defined by the following expression:

 C  D  K eq  a b  A   B c

d

where [A] is the molar concentration of species A at equilibrium, and so forth. The coefficients a, b, c, and d in the chemical equation become exponents in the expression for Keq. Easy 73. Explain the importance of the equilibrium constant.

The Keq is a characteristic numerical value for a given reaction at a given temperature; that is, each chemical reaction has its own characteristic Keq. The concentration of each reactant and product in a chemical reaction at equilibrium is related; the concentrations cannot be random values, but they depend on each other. The numerator of the expression for Keq has the concentrations of every product (however many products there are), while the denominator of the expression for Keq has the concentrations of every reactant, leading to the common products over reactants definition for the Keq. Moderate 74. How do we define the equilibrium constant in terms of partial pressures?

For the gas-phase reaction the pressure-based equilibrium constant, KP, is defined as follows:

KP 

PCc  PDd PBb  PCc

where PA is the partial pressure of substance A at equilibrium in atmospheres, and so forth.

Easy 75. Define Le Chatelier‘s principle. If a reaction is exothermic, if the temperature is increased, what which way will the reaction shift?

If an equilibrium is stressed, then the reaction shifts to reduce the stress. If a reaction is exothermic, the heat will be on the product side. Raising the temperature will shift the reaction towards the reactant side.

Moderate 76. How does the addition or removal of a product or a reactant affect the equilibrium of a reaction?

When additional reactant is added, the equilibrium shifts to reduce this stress: it makes more product. When additional product is added, the equilibrium shifts to reactants to reduce the stress. If reactant or product is removed, the equilibrium shifts to make more reactant or product, respectively, to make up for the loss.

Easy 77. How does an equilibrium react to a change in pressure?

Pressure changes do not markedly affect the solid or liquid phases. However, pressure strongly impacts the gas phase. Le Chatelier‘s principle implies that a pressure increase shifts an equilibrium to the side of the reaction with the fewer number of moles of gas,

while a pressure decrease shifts an equilibrium to the side of the reaction with the greater number of moles of gas. If the number of moles of gas is the same on both sides of the reaction, pressure has no effect.

Moderate 78. What is the effect of temperature changes on an equilibrium?

Because temperature is a measure of the energy of the system, increasing temperature can be thought of as adding energy. The reaction will react as if a reactant or a product is being added and will act accordingly by shifting to the other side. For example, if the temperature is increased for an endothermic reaction, essentially a reactant is being added, so the equilibrium shifts toward products. Decreasing the temperature is equivalent to decreasing a reactant (for endothermic reactions) or a product (for exothermic reactions), and the equilibrium shifts accordingly.

Moderate 79. Explain the relationship between Keq and KP. There is a simple relationship between Keq (based on concentration units) and KP (based on pressure units): KP=Keq(RT)Δn where R is the ideal gas law constant (in units of L·atm/mol·K), T is the absolute temperature, and Δn is the change in the number of moles of gas in the balanced chemical equation, defined as ngas,prods − ngas,rcts.

Easy 80. Explain heterogeneous equilibrium with an example. What is the rule for heterogeneous equilibria?

The equilibrium exhibited by an equation that includes more than one phase of matter is called a heterogeneous equilibrium.

The rule for heterogeneous equilibria is as follows: Do not include the concentrations of pure solids and pure liquids in Keq expressions. Only partial pressures for gas-phase substances or concentrations in solutions are included in the expressions of equilibrium constants. As such, the equilibrium constant expression for the above reaction would simply be

K P  PCO2 because the two solids and one liquid would not appear in the expression.

Easy 81. What are the three parts of an ICE chart?

The three parts of an ICE chart are, I  initial concentrations C  change in concentrations E  equilibrium concentrations Easy 82. Set up (but do not solve) an ICE chart for this reaction, given the initial conditions.

The ICE chart is set up like this. First, the initial values: 3N2(g) 2N3(g)

I C E

0.075

0

For 3x of M N2 lost, 2x M of N3 is produced. These expressions go into the change row: 3N2(g) 2N3(g) I 0.075 0 C –3x +2x E At equilibrium, the resulting concentrations will be combination of the initial amount and the changes: 3N2(g) 2N3(g) I 0.075 0 C –3x +2x E 0.075 – 3x +2x Moderate 83. Define an acid dissociation constant with an example.

An acid dissociation constant, Ka, is the equilibrium constant for the dissociation of a weak.

 H+  H 2 PO4–  Ka   H3PO4  Easy 84. Values of Kb are rarely tabulated. Explain Weak bases also have dissociation constants, labeled Kb (the b subscript stands for base).

However, values of Kb are rarely tabulated because there is a simple relationship between the Kb of a base and the Ka of its conjugate acid: Ka × Kb = 1.0 × 10−14 Easy Fill in the Blanks 85. _____ is the point at which the forward and reverse processes balance each other‘s progress.

Chemical equilibrium; Easy 86. The relationship of the amounts of reactants and products at equilibrium is explained by the _____. law of mass action; Easy 87. The term _____means that the forward and reverse chemical reactions still occur; there is just no net change in the position of the reaction. dynamic; Easy

88. The _____ is a numerical value that relates to the ratio of products and reactants at equilibrium. equilibrium constant; Easy 89. The relationship between Keq and KP is shown by the expression_____.

KP = Keq(RT)Δn; Easy 90. If a chemical equilibrium reaction includes more than one phase of matter, it is said to be in a state of _____. heterogeneous equilibrium; Easy 91. A(n)_____is a change in a condition of the equilibrium. stress; Easy 92. _____ states that if equilibrium is stressed, then the reaction shifts to reduce the stress.

Le Chatelier’s principle; Easy 93. A(n) _____ is a substance that increases the speed of a reaction. catalyst; Easy 94. The direction of shift in an equilibrium reaction can be predicted for changes in _____, _____, or _____. concentrations, temperature, pressure; Easy 95. The _____ is the sum of the initial and change rows in an ICE chart. equilibrium row; Easy 96. Ka × Kb = _____

1.0 × 10−14; Easy

97. Weak bases have a dissociation constant labeled, ____________.

Kb; Easy 98. The Ksp is a special type of the Keq and applies to compounds that are only _____. slightly soluble; Easy 99. The equilibrium constant for a compound normally considered insoluble is called a(n)_____. solubility constant; Easy Chapter 14 Oxidation and Reduction True/False 1. Atoms in their elemental state are assigned an oxidation number of 0.

True; Easy 2. The oxygen atom in O2 is assigned an oxidation number of –1.

False; Easy

3. The fluorine atom in HF is assigned an oxidation number of +1.

False; Easy 4. The sodium atom in NaOH is assigned an oxidation number of +1.

True; Easy 5. The calcium atom in Ca3(PO4)2 is assigned an oxidation number of +3.

False; Moderate 6. The phosphorous atom in H3PO4 is assigned an oxidation number of +5.

True; Moderate 7. When an oxidation number of an atom is increased in the course of a redox reaction, that atom is being reduced.

False; Easy 8. When an oxidation number of an atom is decreased in the course of a redox reaction, that atom is being oxidized.

False; Easy 9. Given: Na + I2 → NaI The coefficient of Na is 2 in the balanced chemical equation.

True; Easy 10. Given: Ru2S3 + O2 → Ru2O3 + S The coefficient of Ru2S3 is 2 in the balanced chemical equation.

True; Easy 11. Given: Be + 2H+ → Be2+ + H2 The equation is a balanced reaction.

True; Moderate 12. Given: Li + Hg2I2 → 2LiI + 2Hg The equation is a balanced reaction.

False; Moderate Refer to the reaction below for True/False Questions 13-14.

Given: Fe + MnO4−  Fe3+ + Mn 13. If the reaction is taking place in an acidic solution, 6 moles of H2O are required on the product side to balance the equation.

False; Hard 14. If the reaction is taking place in a basic solution the balanced equation of the reaction is 7Fe + 3MnO4−  7Fe3+ + 3Mn + 24OH− + 12H2O.

False; Hard 15. Given: Mo2+ + Br2 → Mo3+ + 2Br− The equation is a balanced reaction.

False; Moderate 16. Given: O2 + 2H2O + I2 → 4OH− + 2I− The equation is a balanced reaction.

False; Moderate Refer to the Table 14.1 for True/ False Questions 17-23.

17. The value of E1/2 for the half reaction Mn2+ + 4H2O → MnO4− + 8H+ + 5e− is 1.51 V. False; Easy 18. The value of E1/2 for the half reaction O2 + 4H+ + 4e− → 2H2O is 1.23 V. True; Easy 19. The voltage of a voltaic cell based on the reaction O2 + 4H+ + 4Cl−→ 2H2O + 2Cl2 is 0.15 V. True; Moderate 20. The reaction taking place in the voltaic cell Ag+ + Fe2+ → Ag + Fe3+ is a nonspontaneous reaction.

False; Moderate 21. The reaction taking place in the voltaic cell Pb2+ + H2 → Pb + 2H+ is a spontaneous reaction.

False; Moderate 22. If the voltage of the voltaic cell based on the reaction Cu2+ + 2I− → I2 + Cu is 0.88 V, the standard reduction potential of the Cu2+ + 2e− → Cu half-cell is 0.54 V. False; Moderate 23. If the voltage of the voltaic cell 6I− + 2NO3− + 8H+ + 6e− → 2NO + 4H2O + 3I2 is 1.50 V, the standard reduction potential of the NO3− + 4H+ + 3e− → NO + 2H2O half-cell reaction is –0.96 V. False; Moderate 24. The half reaction for the electrolysis of elemental chlorine is Cl2+ + 2e− → Cl2.

False; Easy 25. The half reaction for the electrolysis of elemental bromine is 2Br− → Br2 + 2e−.

True; Easy Multiple Choice Questions 26. Redox reactions require _____ to keep track of electrons assigned to each atom in a chemical reaction.

ee. mass numbers ff. oxidation numbers gg. number of neutrons hh. magnetic quantum numbers ii. spin quantum numbers b; Easy 27. Atoms in their elemental state are assigned an oxidation number of _____. a. 1 b. 2 c. –2 d. 3 e. 0 e; Easy 28. The oxidation number for the chlorine atom in Cl2 is _____. a. 0 b. +3 c. +1 d. +2 e. –2 a; Easy 29. The oxidation number for the fluorine atom in CaF2 is ____. a. 0 b. –3 c. –1 d. +2 e. –2 c; Easy 30. The oxidation number for the oxygen atom in BaO2 is ____. a. –3 b. –1 c. –2 d. +2 e. 0 b; Moderate 31. The oxidation number for the calcium atom in Ca3(PO4)2 is _____. a. +4 b. +3 c. –3 d. +2 e. +6 d; Moderate 32. The oxidation number of the hydrogen atom in InH3 is _____. a. +2 b. –2 c. –1

d. 0 e. –3 c; Moderate Refer to the reaction below for Multiple Choice Questions 33-34.

Given: 2RbOH + H2SO4  Rb2SO4 + 2H2O 33. What happens to the oxidation number of the Rb atom? a. The oxidation number remains the same at –2. b. The oxidation number changes from 0 to +1. c. The oxidation number remains the same at +1. d. The oxidation number changes from –2 to +1. e. The oxidation number changes from –1 to +2. c; Moderate 34. What happens to the oxidation number of the S atom? a. The oxidation number remains the same at +2. b. The oxidation number changes from 0 to +2. c. The oxidation number remains the same at +6. d. The oxidation number changes from –2 to +2. e. The oxidation number remains the same at +4. c; Moderate 35. Given: Fe + I2 → FeI2 Identify what is being oxidized and reduced in this redox reaction. a. Oxidation numbers of all the atoms remain the same. b. Fe is being reduced from 0 to –2; I is being oxidized from –1 to 0. c. Fe is being oxidized from 0 to +1; I is being reduced from +2 to 0. d. Fe is being reduced from 0 to –1; I is being oxidized from 0 to +2. e. Fe is being oxidized from 0 to +2; I is being reduced from 0 to –1. e; Moderate 36. Given: NH3 + HCl  NH4Cl Which of the following is taking place in the reaction? a. The oxidation number of the hydrogen atom increases from +3 to +4; so it is being oxidized. b. The oxidation number of the chlorine atom decreases from 0 to +1; so it is being oxidized. c. The oxidation number of the chlorine atom decreases from +1 to 0; so it is being reduced. d. The oxidation number of the nitrogen atom increases from –3 to +4, so it is being oxidized. e. The oxidation number of the nitrogen atom remains the same at –3. e; Moderate 37. Given: 6H++ 2MnO4− + 5H2O2 → 2Mn+2 + 5O2 + 8H2O

Identify what is being oxidized and reduced in this redox reaction. a. Oxidation numbers of all the atoms remain the same. b. H is being oxidized from +6 to +8; Mn stays the same. c. O is being oxidized from –1 to 0; Mn is being reduced from +7 to +2. d. Mn is being oxidized from –4 to –2; O is being reduced from +6 to +5. e. H is being oxidized from +6 to +5; Mn is being reduced from +1 to –2. c; Moderate 38. Given: HCN + NaOH  NaCN + H2O Which of the following is taking place in the reaction? a. The oxidation numbers of all the atoms remain the same. b. The oxidation number of the nitrogen atom increases from +3 to +4. c. The oxidation number of the nitrogen atom decreases from +4 to +3. d. The oxidation number of the hydrogen atom decreases from +2 to +1. e. The oxidation number of the carbon atom increases from +2 to +3. a; Moderate Refer to the reaction below for Multiple Choice Questions 39-40.

Given: 2CH4 + 3O2  2CO + 4H2O 39. C is being _____. a. oxidized from +3 to +4 b. oxidized from –4 to +2 c. reduced from +4 to –2 d. reduced from –1 to –2 e. reduced from –1 to –3 b; Moderate 40. What happens to the H atom? a. Oxidation number increases from –1 to +1 b. Oxidation number decreases from +1 to 0 c. Oxidation number remains the same at +1 d. Oxidation number decreases from –1 to 0 e. Oxidation number remains the same at +2 c; Moderate Refer to the reaction below for Multiple Choice Questions 41-43.

Given: Mg + O2  MgO 41. The coefficient of Mg after balancing the equation is _____. a. 1 b. 2 c. 5 d. 3 e. 4 b; Easy 42. The coefficient of O2 after balancing the equation is _____.

a. 1 b. 2 c. 5 d. 3 e. 4 a; Easy 43. The coefficient of MgO after balancing the equation is _____. a. 1 b. 2 c. 5 d. 3 e. 4 b; Easy Refer to the reaction below for Multiple Choice Questions 44-47.

Given: P2O5 + I2 → PI3 + O2 44. The coefficient of P2O5 after balancing the equation is _____. a. 3 b. 5 c. 6 d. 2 e. 4 d; Moderate 45. The coefficient of I2 after balancing the equation is _____. a. 3 b. 5 c. 6 d. 2 e. 4 c; Easy 46. The coefficient of PI3 after balancing the equation is _____. a. 3 b. 5 c. 6 d. 2 e. 4 e; Easy 47. The coefficient of O2 after balancing the equation is _____. a. 3 b. 5 c. 6 d. 2 e. 4 b; Easy

Refer to the reaction below for Multiple Choice Questions 48-51.

Given: Sr + H+→ Sr2+ + H2 48. The coefficient of Sr after balancing the equation is _____. a. 1 b. 5 c. 3 d. 2 e. 4 a; Easy 49. The coefficient of H+ after balancing the equation is _____. a. 1 b. 5 c. 3 d. 2 e. 4 d; Easy 50. The coefficient of Sr2+ after balancing the equation is _____. a. 1 b. 5 c. 3 d. 2 e. 4 a; Easy 51. The coefficient of H2 after balancing the equation is _____. a. 1 b. 5 c. 3 d. 2 e. 4 a; Easy Refer to the reaction below for Multiple Choice Questions 52-54.

Given: Au+ → Au + Au3+ (Hint: both half reactions will start with the same reactant.)

52. The coefficient of Au+ after balancing the equation is _____. a. 1 b. 5 c. 3 d. 2 e. 4 c; Moderate 53. The coefficient of Au after balancing the equation is _____. a. 1 b. 5

c. 3 d. 2 e. 4 d; Moderate 54. The coefficient of Au3+ after balancing the equation is _____. a. 1 b. 5 c. 3 d. 2 e. 4 a; Moderate Refer to the reaction below for Multiple Choice Questions 55-58.

Given: Cl− + F2 → F− + Cl2 55. The coefficient of Cl− after balancing the equation is _____. a. 1 b. 5 c. 3 d. 2 e. 4 d; Moderate 56. The coefficient of F2 after balancing the equation is _____. a. 1 b. 5 c. 3 d. 2 e. 4 a; Moderate 57. The coefficient of Cl2 after balancing the equation is _____. a. 2 b. 5 c. 3 d. 1 e. 4 d; Moderate 58. The coefficient of F− after balancing the equation is _____. a. 1 b. 2 c. 5 d. 3 e. 4 b; Moderate Refer to the reaction below for Multiple Choice Questions 59-62.

Given:

Cr2O72− + Fe → Cr3+ + Fe3+ 59. If the reaction is taking place in a basic solution, what is the number of moles of water required on the product or reactant side to balance the redox reaction? a. 7 moles of H2O on the reactant side b. 7 moles of H2O on the product side c. 14 moles of H2O on the reactant side d. No moles of H2O e. 14 moles of H2O on the product side a; Hard 60. If the reaction is taking place in an acidic solution, what is the number of moles of water required on the product or reactant side to balance the redox reaction? a. 14 moles of H2O on the reactant side b. 14 mole2 of H2O on the product side c. No moles of H2O d. 7 moles of H2O on the reactant side e. 7 moles of H2O on the product side e; Hard 61. If the reaction is taking place in an acidic solution, the balanced equation is: a. 2Fe + 14H+ + 2Cr2O72− → 2Fe3+ + 2Cr3+ + 2H2O. b. Fe + H+ + Cr2O72− → Fe3+ + Cr3+ + H2O. c. 2Fe + 14H+ + Cr2O72− → 4Fe3+ + 2Cr3+ + 28H2O. d. 2Fe + 14H+ + Cr2O72− → 2Fe3+ + 2Cr3+ + 7H2O. e. 4Fe + 10H+ + Cr2O72− → 3Fe3+ + Cr3+ + 7H2O. d; Hard 62. If the reaction is taking place in a basic solution, the balanced equation is: a. 14H2O + 2Cr2O72− + 2Fe  2Cr3+ + 14OH− + 4Fe3+. b. 14H2O + Cr2O72− + 5Fe  2Cr3+ + 14OH− + 5Fe3+. c. 7H2O + Cr2O72− + 2Fe  2Cr3+ + 14OH− + 2Fe3+. d. H2O + Cr2O72− + Fe  Cr3+ + OH− + Fe3+. e. 7H2O + 7Cr2O72− + 7Fe  7Cr3+ + 14OH− + 7Fe3+. c; Hard Refer to the Table 14.1 for Multiple Choice Questions 63-67.

63. What is the E of a voltaic cell based on this reaction?

F2 + Pb → 2F– + Pb2+ a. 2.74 V b. 22.7 V c. 2.99 V d. 0.362 V e. 5.49 V c; Easy 64. What is the E of a voltaic cell based on this reaction?

2Ce4+ + Ni → 2Ce3+ + Ni2+ a. 2.36 V

b. 2.11 V c. 6.44 V d. 0.403 V e. 1.86 V e; Easy 65. What is the E of a voltaic cell based on this reaction?

2Mn2+ + 8H2O + 5Sn4+ → 2MnO4– + 16H+ + 5Sn2+ a. –10.1 V b. –1.36 V c. +10.1 V d. +0.227 V e. +1.81 V b; Moderate 66. What is the E of a voltaic cell based on this reaction?

NO3− + 4H+ + Fe2+ → NO + 2H2O + Fe3+ a. +0.21 V b. +0.77 V c. –0.77 V d. –0.73 V e. +0.35 V a; Moderate 67. What is the E of a voltaic cell based on this reaction?

MnO4− + 6H+ + H2 → Mn2+ + 4H2O a. –1.10 V b. +1.10 V c. +1.51 V d. +3.68 V e. +1.23 V c; Moderate 68. What is the half reaction for the electrolysis of elemental nickel? a. Ni2+ + 2e− → Ni b. Ni+ + e− → Ni c. Ni3++ 3e− →Ni d. Ni2 + 2e− → 2Ni− e. 2Ni+ + 2e− → Ni2 a; Easy 69. What is the half reaction for the electrolysis of elemental fluorine? a. F2+ + 2e− → F b. F+ + e− → F c. F3++ 3e− →F d. F2 → 2F− + 2e− e.

2F−  F2 + 2e− e; Easy 70. What is the half reaction for the electrolysis of elemental iodine? a. I2 → 2I− + 2e−

b. I2+ + 2e− → I c. 2I− → I2 + 2e− d. 2I+ + 2e− → I2 e. I2 + 2e− → 2I− c; Easy Essay Questions 71. Define oxidation, reduction, and oxidation-reduction reactions.

Oxidation is defined as the loss of one or more electrons by an atom.

Reduction is defined as the gain of one or more electrons by an atom.

Chemical reactions that involve the transfer of electrons are called oxidationreduction reactions.

Easy 72. Briefly describe the rules required for assigning oxidation numbers.

The rules for assigning oxidation numbers to atoms are as follows: 1. Atoms in their elemental state are assigned an oxidation number of 0.

2. Atoms in monatomic (i.e., one-atom) ions are assigned an oxidation number equal to their charge. Oxidation numbers are usually written with the sign first, then the magnitude, to differentiate them from charges.

3. In compounds, fluorine is assigned a −1 oxidation number; oxygen is usually assigned a −2 oxidation number [except in peroxide compounds (where it is −1) and in binary compounds with fluorine (where it is positive)]; and hydrogen is usually assigned a +1 oxidation number [except when it exists as the hydride ion (H−), in which case rule 2 prevails].

4. In compounds, all other atoms are assigned an oxidation number so that the sum of the oxidation numbers on all the atoms in the species equals the charge on the species (which is zero if the species is neutral).

Moderate 73. Briefly describe the method of balancing a redox reaction by inspection.

Balancing simple redox reactions can be a straightforward matter of going back and forth between products and reactants. For example, in the redox reaction of Na and Cl2: Na + Cl2 → NaCl it should be immediately clear that the Cl atoms are not balanced. We can fix this by putting the coefficient 2 in front of the product: Na + Cl2 → 2NaCl However, now the sodium is unbalanced. This can be fixed by including the coefficient 2 in front of the Na reactant: 2Na + Cl2 → 2NaCl We say that we are able to balance the reaction by inspection. Many simple redox reactions can be balanced by inspection.

Easy 74. Briefly describe the method of balancing a redox reaction by the half reaction method.

Consider this redox reaction: Al + Ag+ → Al3+ + Ag

To balance this, we will write each oxidation and reduction reaction separately, listing the number of electrons explicitly in each. Individually, the oxidation and reduction reactions are called half reactions.

We will then take multiples of each reaction until the number of electrons on each side cancels completely and combine the half reactions into an overall reaction, which should then be balanced. This method of balancing redox reactions is called the half reaction method.

This half reaction is not completely balanced because the overall charges on each side are not equal. When an Al atom is oxidized to Al3+, it loses three electrons. We can write these electrons explicitly as products: Al → Al3+ + 3e− Now this half reaction is balanced—in terms of both atoms and charges.

The reduction half reaction involves silver: Ag+ → Ag The overall charge is not balanced on both sides. But we can fix this by adding one electron to the reactant side because the Ag+ ion must accept one electron to become the neutral Ag atom: Ag+ + e− → Ag This half reaction is now also balanced.

When combining the two half reactions into a balanced chemical equation, the key is that the total number of electrons must cancel, so the number of electrons lost by atoms are equal to the number of electrons gained by other atoms.

With three electrons as products and one as reactant, the least common multiple of these two numbers is three: we can use a single aluminum reaction but must take three times the silver reaction: Al → Al3+ + 3e− 3 × [Ag+ + e− → Ag] The 3 on the second reaction is distributed to all species in the reaction: Al → Al3+ + 3e− 3Ag+ + 3e− → 3Ag Now the two half reactions can be combined just like two algebraic equations, with the arrow serving as the equals sign. The same species on opposite sides of the arrow can be canceled: Al + 3Ag++ 3e−→Al3++ 3Ag + 3e− The net balanced redox reaction is as follows: Al + 3Ag+ → Al3+ + 3Ag Moderate 75. Briefly enumerate the steps involved in balancing a redox reaction with an example.

Assume a basic solution.

Assume the reaction: MnO2 + CrO3− → Mn + CrO4− We start by separating the oxidation and reduction processes so we can balance each half reaction separately. The oxidation reaction is as follows: CrO3− → CrO4− The Cr atom is going from a +5 to a +7 oxidation state and loses two electrons in the process. We add those two electrons to the product side: CrO3− → CrO4− + 2e− Now we must balance the O atoms. Because the solution is basic, we should use OH− rather than H2O:

OH− + CrO3− → CrO4− + 2e− We have introduced H atoms as part of the reactants; we can balance them by adding H+ as products: OH− + CrO3− → CrO4− + 2e− + H+ If we check the atoms and the overall charge on both sides, we see that this reaction is balanced. However, if the reaction is occurring in a basic solution, it is unlikely that H+ ions will be present in quantity. The way to address this is to add an additional OH− ion to each side of the equation: OH− + CrO3− + OH− → CrO4− + 2e− + H+ + OH− The two OH− ions on the left side can be grouped together as 2OH−. On the right side, the H+ and OH− ions can be grouped into an H2O molecule: 2OH− + CrO3− → CrO4− + 2e− + H2O This is a more appropriate form for a basic solution.

Now we balance the reduction reaction: MnO2 → Mn The Mn atom is going from +4 to 0 in oxidation number, which requires a gain of four electrons: 4e− + MnO2 → Mn Then we balance the O atoms and then the H atoms: 4e− + MnO2 → Mn + 2OH− 2H+ + 4e− + MnO2 → Mn + 2OH− We add two OH− ions to each side to eliminate the H+ ion in the reactants; the reactant species combine to make two water molecules, and the number of OH− ions in the product increases to four: 2H2O + 4e− + MnO2 → Mn + 4OH− This reaction is balanced for a basic solution.

Now we combine the two balanced half reactions. The oxidation reaction has two electrons, while the reduction reaction has four. The least common multiple of these two numbers is four, so we multiply the oxidation reaction by 2 so that the electrons are balanced: 2 × [2OH− + CrO3− → CrO4− + 2e− + H2O] 2H2O + 4e− + MnO2 → Mn + 4OH− Combining these two equations results in the following equation: 4OH− + 2CrO3− + 2H2O + 4e− + MnO2 → 2CrO4− + 4e− + 2H2O + Mn + 4OH− The four electrons cancel. So do the two H2O molecules and the four OH− ions. What remains is 2CrO3− + MnO2 → 2CrO4− + Mn which is our final balanced redox reaction.

Moderate 76. Describe a voltaic cell.

Consider this redox reaction: Zn + Cu2+ → Zn2+ + Cu If you were to mix zinc metal and copper ions in a container, this reaction would proceed by itself; we say that this reaction is spontaneous. Suppose, however, we set up this reaction in a way depicted in the figure below. Zinc and zinc ions are on one side of the system, while copper and copper ions are on the other side of the system. The two parts are connected with a wire.

Even though the two half reactions are physically separated, a spontaneous redox reaction still occurs. However, in this case, the electrons transfer through the wire connecting the two half reactions; that is, this setup becomes a source of electricity. Useful work can be extracted from the electrons as they transfer from one side to the other—for example, a light bulb can be lit, or a motor can be operated. The apparatus as a whole, which allows useful electrical work to be extracted from a redox reaction, is called a voltaic (galvanic) cell.

Moderate 77. Define half cell, anode, cathode, and electrodes.

Each individual system that contains a half reaction is called a half cell.

The half cell that contains the oxidation reaction is called the anode.

The half cell that contains the reduction reaction is called the cathode.

The cathode and anode collectively are the electrodes of the voltaic cell.

Easy 78. Define salt bridge, voltage, and standard reduction potentials.

A salt bridge contains a solution of some ionic compound whose ions migrate to either side of the voltaic cell to maintain the charge balance.

The tendency for electrons to go from one half cell to another is called the voltage of the voltaic cell, represented by E. Standard reduction potential is the voltage of a reduction half reaction relative to the hydrogen half reaction.

Easy 79. Describe the dry cell.

In 1866, the French scientist Georges Leclanché invented the dry cell. A schematic of a dry cell is shown in the figure below.

The zinc case and the central carbon rod serve as the anode and cathode, respectively.

The other reactants are combined into a moist paste that minimizes free liquid, so the battery is less messy (hence the name dry cell). The actual redox reaction is complex but can be represented by the following redox reaction: Zn + 2MnO2 + 2NH4+ → Zn2+ + Mn2O3 + 2NH3 + H2O A dry cell has a voltage of about 1.56 V. While common and useful, dry cells have relatively short lifetimes and contain acidic components. They also cannot be recharged, so they are one-use only. Batteries that can be used only once are called primary batteries.

Moderate 80. Describe the alkaline battery.

In the late 1950s, Lewis Urry of the Eveready Battery Company in Ohio invented the alkaline battery. Alkaline batteries are similar to dry cells, but they use a basic moist paste rather than an acidic one. Moreover, the net amount of base does not change during the course of the redox reaction. The overall redox reaction is as follows: Zn + 2MnO2 → ZnO + Mn2O3 Alkaline batteries have the advantage of being longer lasting and holding their voltage better—about 1.54 V—throughout their lifetime.

Easy 81. Describe button batteries.

A common type of battery, especially with the increased popularity of personal electronic devices, is the button battery. A button battery is a small battery that can power small electronic devices; the batteries can be as small as 5 mm across. Two popular redox reactions used for button batteries are the alkaline dry-cell reaction and a silver oxide-based reaction: Zn + Ag2O → ZnO + 2Ag

Some button batteries use a lithium-based redox reaction, typified by this anode reaction: Li → Li+ + e− E1/2 = 3.045 V Moderate 82. Describe a lead storage battery.

An important secondary battery is the lead storage battery. The lead storage battery is based on this redox reaction: Pb + PbO2 + 4H+ + SO42− → 2PbSO4 + 2H2O The redox reaction produces about 2 V, but it is typical to tie several individual batteries together to generate a larger voltage. The lead storage battery has the distinction that the product of both half reactions is PbSO4, which as a solid accumulates on the many plates within each cell. The lead storage battery is a secondary battery, as it can be recharged and reused many times. Because it is based on lead, these batteries are rather heavy. They should also be recycled when replaced so that potentially dangerous lead does not escape into the environment. Because of their characteristics, lead storage batteries are used to start large engines in automobiles, boats, and airplanes.

Moderate 83. Define an electrolytic cell and electrolysis.

By forcing electricity into a voltaic cell, we can make a redox reaction occur that normally would not be spontaneous. Under these circumstances, the cell is called an electrolytic cell, and the process that occurs in the cell is called electrolysis.

Easy 84. Explain how elements are isolated using electrolysis with an example.

If NaCl is melted at about 800°C in an electrolytic cell and an electric current is passed through it, elemental sodium will appear at the cathode and elemental chlorine will appear at the anode as the following two reactions occur: Na+ + e− → Na 2Cl− → Cl2 + 2e− Normally we expect elemental sodium and chlorine to react spontaneously to make NaCl. However, by using an input of electricity, we can force the opposite reaction to occur and generate the elements. Lithium, potassium, and magnesium can also be isolated from compounds by electrolysis.

Easy 85. Describe electroplating.

Electroplating is the deposition of a thin layer of metal on an object for protective or decorative purposes. Essentially, a metal object is connected to the cathode of an electrolytic cell and immersed in a solution of a particular metal cation. When the electrolytic cell is operated, a thin coating of the metal cation is reduced to the elemental metal on the surface of the object; the thickness of the coating can be as little as a few micrometers (10−6 m). Jewelry, eating utensils, electrical contacts, and car parts like bumpers are common items that are electroplated. Gold, silver, nickel, copper, and chromium are common metals used in electroplating.

Easy Fill in the Blanks 86. Consider the chemical reaction:

2Li(s) + I2(g)  2LiI(s) The reactants are two electrically neutral elements; they have the same number of electrons as protons.

This reaction involves the _____between atoms. transfer of electrons; Easy 87. _____is defined as the loss of one or more electrons by an atom. oxidation; Easy 88. _____is defined as the gain of one or more electrons by an atom. reduction; Easy 89. Redox reactions use _____ to keep track of the electrons assigned to each atom in a chemical reaction. oxidation numbers; Easy 90. Chemical reactions that involve the transfer of electrons are called _____. oxidation-reduction (or redox reactions); Easy 91. Given: K + Br2 → KBr Coefficient of K in the balanced equation is____.

2; Easy 92. Given: B2O3 + H2 → B + H2O Coefficient of B2O3 in the balanced equation is ____.

1; Moderate 93. An apparatus as a whole, which allows useful electrical work to be extracted from a redox reaction, is called a(n)_____. voltaic (galvanic) cell; Easy 94. Each individual system that contains a half reaction is called a(n)_____. half cell; Easy 95. The half cell that contains the oxidation reaction is called the_____. anode; Easy 96. The half cell that contains the reduction reaction is called the_____. cathode; Easy 97. The cathode and anode collectively are the _____ of the voltaic cell. electrodes; Easy 98. The tendency for electrons to go from one half cell to another is called the _____. voltage; Easy 99. The voltage of a reduction half reaction relative to the hydrogen half reaction is called the ______. standard reduction potential; Easy

100. A cell into which electricity is forced to make a nonspontaneous reaction occur is called a(n) _____. electrolytic cell; Easy Chapter 15 Nuclear Chemistry True/False 1. A 115 B atom has 5 protons in its nucleus.

True; Easy 2. An 27 13 Al atom has13 neutrons in its nucleus.

False; Easy 3. An alpha particle is a collection of four protons and five neutrons and is equivalent to a beryllium nucleus.

False; Easy 4. A beta particle is a neutron emitted from the nucleus.

False; Easy 5. Gamma rays are high-energy electromagnetic radiation given off in radioactive decay.

True; Easy 6. Alpha particles penetrate more than beta particles.

False; Moderate 7. Nuclear fusion is the breaking down of large nuclei into smaller nuclei, usually with the release of excess neutrons.

False; Moderate 8. The reactant in the nuclear equation is the parent isotope.

True; Easy 9. The product in the nuclear equation is the daughter isotope.

True; Easy 10. Only radioactive isotopes have a half-life.

True; Moderate 11. An isotope with a longer half-life is generally considered more radioactive.

False; Moderate 12. Radioactive decay is a linear process.

False; Easy 13. 1 Gy = 1,000 rad False; Easy

14. rem = 1 Ci  factor False; Easy 15. 1 rad = 1 J of energy per gram of tissue False; Easy 16. 1 Bq = 10 decays/s False; Easy 17. 1 Bq = 1 Ci False; Easy 18. Beta particles contribute more to the rems of exposure than alpha particles.

False; Moderate 19. Cobalt-60 is used in irradiation of food.

True; Moderate 20. 99mTc is used in intestinal scanning.

True; Easy 21. 133Xe is used in lung imaging.

True; Easy 22. 59Fe is used in cancer detection.

False; Easy 23.

According to Albert Einstein‘s theory of relativity, energy (E) and mass (m) are related by the following equation: E = mc where c is the speed of light, or 3.00 × 108 m/s.

False; Easy 24. Uranium-238 can absorb a neutron and undergo a fission reaction to produce an atom of cesium-135 and an atom of rubidium-96. The balanced nuclear equation for the process is 238 92

135 1 U + 10 n 96 37 Rb + 55 Cs + 80 n.

True; Moderate 25. A chain reaction is a process that generates more reaction pathways for each previous reaction.

True; Moderate Multiple Choice Questions 26. Complete the reaction.

235 92

U  24 He  90? Th

z. 230 aa. 231 bb. 232 cc. 233 dd. 234 b; Easy

27. Francium has an atomic number of 87. Which of the following is the nuclear equation that represents the radioactive decay of francium-223 by alpha particle?

2 221 a. 223 87 Fr  1 H  86 Rn b.

223 87

Fr  49 Be  214 83 Bi

c.

223 87

Fr  126 C  21181Th

d.

223 87 223 87

Fr  24 He  219 85 At

e.

Fr  73 Li  216 84 Po d; Easy 28. Polonium has an atomic number of 84. When polonium-209 emits an alpha particle, the resulting daughter isotope is _____. a. 205 85 At b.

205 86

Rn

c.

205 84

Po

d.

205 83 205 82

Bi

e.

Pb e; Easy 29. Actinium has an atomic number of 89. When _____ emits an alpha particle, actinium-227 is the resulting daughter isotope. a. 231 86 Rn b. c. d.

231 87 231 88 231 93

Fr Ra

Np

e. 231 91 Pa e; Easy 30. Dubnium has an atomic number of 105. When dubnium-262 emits an alpha particle, _____ is the resulting daughter isotope. a. 258 101 Md b. c. d.

258 103 258 88

Lr

258 102 258 85

No

Ra

e.

At b; Easy 31. Neptunium has an atomic number of 93. When _____ emits an alpha particle, neptunium239 is the resulting daughter isotope. a. 243 91 Pa b.

243 88

Ac

c. d.

243 92 243 95

U Am

243 94

e.

Pu d; Easy 32. Complete the following reaction.

12 6

C  127 N  ? 

c.

0 1 4 2 0 1

d.

1 0

a. b.

e

He

p n

e. H a; Moderate 33. Radium has an atomic number of 88. _____is the daughter isotope formed when radium226 emits a beta particle. a. 226 88 Ra b. c. d.

226 86 226 89

Rn

Ac

226 87 226 90

Fr

e.

Th c; Moderate 34. Lawrencium has an atomic number of 103. _____is the daughter isotope formed when lawrencium-262 emits a beta particle. a. 262 102 No b. c. d.

262 107 262 106 262 104 262 105

Bh Sg Rf

e.

Db d; Moderate 35. Radium has an atomic number of 88. Radium-226 is the daughter isotope formed when _____ emits a beta particle. a. 262 89 Ac b.

226 87

Fr

c.

226 88 226 86 226 78

Ra

d. e.

Rn Pt

b; Moderate 36. Uranium has an atomic number of 92. Uranium-238 is the daughter isotope formed when _____ emits a beta particle. a. 238 94 Pu b. c. d.

238 93 238 92

Np U

238 91 238 90

Pa

e.

Th d; Moderate 37. Energies of gamma rays are typically expressed in units of megaelectron volts (MeV), where 1 MeV = 1.602 × 10−13 J. Energies of gamma rays emitted when oxygen-13 gives off a beta particle is 0.168 MeV. What is its energy in joules? a. 6.29 × 10–12J b. 5.01 × 1012J c. 1.05 × 10–12J d. 0.51 × 1014 J e. 2.69 × 10 –14 J e; Moderate 38. Energies of gamma rays are typically expressed in units of megaelectron volts (MeV), where 1 MeV = 1.602 × 10−13 J. Energies of gamma rays emitted when nitrogen-11 gives off a beta particle is 1.508 × 10-12 J. What is its energy in MeV? a. 9.14 MeV b. 4.19 MeV c. 9.41 MeV d. 1.49 MeV e. 4.91 MeV c; Moderate 39. How long does it take for 10.0 g of a radioactive isotope to decay to 1.25 g if its half-life is 17.0 d? a. 15.0 d b. 51.0 d c. 17.0 d d. 71.0 d e. 30.0 d b; Moderate 40. The half-life of a radioactive sample is 11.0 s. If the sample initially contains 25.0 g of the radioactive sample, how much remains after 54.0 s? a. 0.00338 g b. 0.0338 g c. 0.833 g d. 0.383 g e. 0.338 g c; Moderate

41. The half-life of a radioactive sample is 60.0 y. If the sample weighs 0.893 g initially, how much remains after 420.0 y? a. 0.00698 g b. 0.00968 g c. 0.00986 g d. 0.00689 g e. 0.00896 g a; Moderate 42. How long does it take for 5.05 g of a radioactive isotope to decay to 0.0505 g if its halflife is 20,000 y? a. 205,000 y b. 505,000 y c. 313,000 y d. 331,000 y e. 133,000 y e; Moderate 43. It took 65.5 y for 30.0 g of a radioactive isotope to decay to 1.25 g. What is the half-life of this isotope? a. 14.3 y b. 3.14 y c. 4.13 y d. 13.4 y e. 31.4 y a; Moderate 44. It took 70.2 s for 4.86 g of a radioactive isotope to decay to 0.0256 g. What is the half-life of this isotope? a. 7.29 s b. 4.19 s c. 9.27 s d. 7.92 s e. 14.9 s c; Moderate 45. The half-life of americium-241 is 432 y. If 2.00 g of americium-241 is present in a sample, what mass of americium-241 is present after 1,000.0 y? a. 0.0402 g b. 0.0204 g c. 0.204 g d. 0.402 g e. 0.420 g d; Moderate 46. If the half-life of tritium (hydrogen-3) is 12.3 y, how much of a 0.0666 g sample of tritium is present after 50.0 y? a. 8.93  10–3 g b. 9.83  10–3 g c. 3.98  10–3 g

d. 8.39  10–3 g e. 8.93  10–3 g c; Moderate 47. A sample containing carbon-14 contains 5.30 × 10−6 g of carbon-14 in it. If the age of the sample is 15,800 y, how much carbon-14 did it have originally? The half-life of carbon14 is 5,730 y. a. 3.58  10–6 g b. 3.58  10–5 g c. 3.58  10–4 g d. 8.53  10–5 g e. 8.53  10–4 g b; Moderate 48. The half-life of carbon-11 is 20.3 min. If 5.00 g of carbon-11 is left in the sample after 59.3 min, what mass of carbon-11 was present initially? a. 93.7 g b. 79.3 g c. 73.9 g d. 37.9 g e. 39.7 g d; Moderate 49. A sample of radon has an activity of 90,000 Bq. If the half-life of radon is 15 h, how long before the sample‘s activity is 5,625 Bq? a. 60 h b. 90 h c. 15 h d. 30 h e. 45 h a; Easy 50. A sample of radon has an activity of 80,000 Bq. If the half-life of radon is 15 h, how long before the sample‘s activity is 5,000 Bq? a. 10 h b. 20 h c. 45 h d. 30 h e. 60 h e; Easy 51. _____ = 3.7  1010 decays/s a. 1 rad b. 1 rem c. 1 Ci d. 1 cal e. 1 coulomb c; Easy

52. A sample of radon gas has an activity of 140.0 mCi. If the half-life of radon is 1,500.0 y, how long before the activity of the sample is 10.0 mCi? a. 1,750 y b. 1,510 y c. 7,510 y d. 5,170 y e. 5,710 y e; Moderate 53. A sample of curium has an activity of 2,450 Bq. If the half-life of curium is 24.0 s, how long before its activity is 25.0 Bq? a. 591 s b. 195 s c. 159 s d. 951 s e. 915 s c; Moderate 54. If a radioactive sample has an activity of 65.0 nCi, how many disintegrations per second are occurring? a. 1,240 disintegrations/s b. 2,140 disintegrations/s c. 2,410 disintegrations/s d. 1,240 disintegrations/s e. 1,420 disintegrations/s c; Moderate 55. If a radioactive sample has an activity of 8.33 Bq, how many disintegrations per second are occurring? a. 8.33 disintegrations/s b. 4.16 disintegrations/s c. 3.38 disintegrations/s d. 30.8 disintegrations/s e. 33.8 disintegrations/s a; Moderate 56. A sample of fluorine-20 has an activity of 5.94 mCi. If its half-life is 11.0 s, what is its activity after 55.0 s? a. 0.816 mCi b. 0.861 mCi c. 0.168 mCi d. 0.186 mCi e. 0.618 mCi d; Moderate 57. A radioactive sample has a half-life of 28.1 y. If 98.0 Bq of the sample were allowed to decay for 27.0 y, what would the activity of the remaining sample be? a. 53.7 Bq b. 537 Bq c. 0.735 Bq d. 7.35 Bq

e. 50.4 Bq e; Moderate 58. How long does it take 200.0 mCi of a radioactive sample to decay to 20.0 mCi if its halflife is 11.0 s? a. 25.4 s b. 5.36 s c. 36.5 s d. 53.6 s e. 73.0 s c; Moderate 59. A typical dose of a radioactive sample is 27.0 mCi. How long does it take for the activity to reduce to 0.100 mCi? The half-life of the sample is 211,000 y. a. 0.17  105 y b. 0.17  106 y c. 1.70  105 y d. 1.70  104 y e. 1.70  106 y e; Moderate 60. After chemical analysis, a radioactive sample is found to contain 2 g of uranium to every 5 g of thorium, its daughter isotope. If the half-life of uranium is 68.9 y, approximately how old is the sample? a. 70.0 y b. 83.7 y c. 111 y d. 165 y e. 125 y e; Moderate 61. After chemical analysis, a radioactive sample is found to contain 1.00 g of francium to every 2.50 g of astatine, its daughter isotope. If the half-life of francium is 4.18 min, approximately how old is the sample? a. 7.65 min b. 7.56 min c. 5.67 min d. 5.76 min e. 6.57 min b; Moderate 62. Complete the following reaction.

226

Ra 156 Pm  68 Co  ?1n

a. 0 b. 1 c. 4 d. 2 e. 3 d; Moderate

63. Complete the following reaction.

235

U + 1 n 141 Ba  92 Kr  ?1n

a. 0 b. 1 c. 4 d. 2 e. 3 e; Moderate 64. Complete the following reaction.

235

U + 1 n 152 Nd  81 Ge  ?1 n

a. 0 b. 1 c. 4 d. 2 e. 3 e; Moderate 65. For every mole of a radioactive sample that decays, 0.1560 g of mass is lost. How much energy is given off per mole of radioactive sample reacted? a. 1.04  1013 J b. 4.01  1013 J c. 4.10  1013 J d. 0.14  1013 J e. 1.40  1013 J e; Moderate 66. For every mole of radioactive sample that decays, 0.2002 g of mass is lost. How much energy is given off per mole of radioactive sample reacted? a. 1.028  1013 J b. 1.208  1013 J c. 1.802  1013 J d. 2.810  1013 J e. 2.180  1013 J c; Moderate 67. Determine the change in mass for the reaction below. Masses in grams are provided.

235

U+

235.0439

1

n  140 Ba  92 Kr  3 1 n

1.0087

139.9106

91.9262

1.0087

a. -0.0087 g b. -1.0087 g c. -1.1897 g d. -0.0439 g e. -0.9262 g c; Moderate 68. Determine the change in mass for the reaction below. Masses in grams are provided.

235

U+

235.0439

1

n  131Sn  103 Mo  2 1 n

1.0087

a. -0.2050 g

130.9170

102.9132

1.0087

b. -0.0439 g c. -0.9170 g d. -0.9132 g e. -0.0087 g a; Moderate 69. What is the energy change of this fission reaction? Masses in grams are provided.

235

U+

235.0439

1

n  140 Ba  92 Kr  3 1 n

1.0087

139.9106

91.9262

1.0087

a. 1.087  1014 J b. 1.043  1014 J c. 1.910  1014 J d. 1.071  1014 J e. 1.926  1014 J d; Moderate 70. Determine the energy change for the reaction below. Masses in grams are provided.

235

U+

235.0439

1

n  131Sn  103 Mo  2 1 n

1.0087

130.9170

102.9132

1.0087

a. 1.044  1013 J b. 1.845  1013 J c. 1.009  1013 J d. 1.917  1013 J e. 1.005  1013 J b; Moderate Essay Questions 71. Define radioactivity.

Radioactivity is the spontaneous emission of particles and electromagnetic radiation from nuclei of unstable atoms.

Easy 72. Describe alpha particle emission.

An alpha particle is composed of two protons and two neutrons and is the same as a helium nucleus. (We often use to represent an alpha particle.) It has a 2+ charge. When a radioactive atom emits an alpha particle, the original atom‘s atomic number decreases by two (because of the loss of two protons), and its mass number decreases by four (because of the loss of four nuclear particles). We can represent the emission of an alpha particle with a chemical equation—for example, the alpha-particle emission of uranium-235 is as follows: 235 92

U  42 He  231 90Th

Easy 73. Describe beta particle emission.

A beta particle is an electron ejected from the nucleus (not from the shells of electrons about the nucleus) and has a 1− charge. We can also represent a beta particle as 01 e . The net effect of beta particle emission on a nucleus is that a neutron is converted to a proton.

The overall mass number stays the same, but because the number of protons increases by one, the atomic number goes up by one. Carbon-14 decays by emitting a beta particle: 14 6

C  147 N  01 e

Easy 74. Describe gamma ray emission.

Gamma rays themselves do not carry an overall electrical charge, but they may knock electrons out of atoms in a sample of matter and make it electrically charged (for which gamma rays are termed ionizing radiation). For example, in the radioactive decay of radon-222, both alpha and gamma radiation are emitted, with the latter having an energy of 8.2  10−14 J per nucleus decayed: 222 86

4 Rn  218 84 Po  2 He  

Easy 75. Define nuclear fission.

Nuclear fission is the breaking down of large nuclei into smaller nuclei, usually with the release of excess neutrons.

Easy 76. Who is the curie named after? Who is the becquerel named after?

The curie is named after Polish scientist Marie Curie, who performed some of the initial investigations into radioactive phenomena in the early 1900s; the becquerel is named after Henri Becquerel, who discovered radioactivity in 1896.

Moderate 77. Give the various expressions used to determine the final amount remaining in a radioactive sample.

We can determine the amount of a radioactive isotope remaining after a given number of half-lives by using the following expression:

1 amount remaining  initial amount    2

n

where n is the number of half-lives. This expression works even if the number of halflives is not a whole number.

If the elapsed time is not an exact number of half-lives, the equation is

amount remaining = (amount initially)  e0.693t / t1/2 where e is the base of natural logarithms (2.71828182…), t is the elapsed time, and t1/2 is the half-life of the radioactive isotope.

Moderate 78. Which is more radioactive—an isotope with a long half-life or an isotope with a short half-life?

An isotope with a shorter half-life is generally considered more radioactive (although it also depends on the energy emitted).

Easy

79. Explain why the amount left in a radioactive sample after 1,000.0 y is not one-tenth of the amount present after 100.0 y, despite the fact that the amount of time elapsed is 10 times as long.

Radioactive decay is an exponential process, not a linear process.

Moderate 80. Define 1 curie (Ci).

The unit curie (Ci) is defined as 3.7  1010 decays/s.

Easy 81. Define 1 becquerel (Bq).

One decay per second is called one becquerel (Bq).

Easy 82. Define 1 rad.

The rad (an acronym for radiation absorbed dose) is a unit equivalent to 1 g of tissue absorbing 0.01 J: 1 rad = 0.01 J/g Easy 83. Plutonium-239 emits alpha particles and is hazardous when inhaled or ingested. What new element is formed by this alpha emission?

Uranium Easy 84. Define 1 rem.

Predicting the effects of radiation is complicated by the fact that different types of emissions affect various tissues differently. To quantify these effects, the unit rem (an acronym for roengten equivalent, man) is defined as rem = rad  factor Easy 85. Explain how a Geiger counter works to detect radiation.

A Geiger counter has a tube with a thin window that allows radioactive emissions through. These radioactive emissions ionize the gas inside the tube, which allows for a current to flow and a ―click‖ to be heard over an audio circuit.

Moderate Fill in the Blanks 86. Three forms of radioactive emissions are ___________ particle, _____________ particle and ________________ rays. alpha, beta, and gamma; Easy 87. ________________ can penetrate deeply into matter and can impart large amounts of energy into surrounding matter.

Gamma rays; Easy 88. A(n) _____ is a chemical equation that emphasizes changes in atomic nuclei. nuclear equation; Easy

89. The breaking apart of an atomic nucleus into smaller nuclei is called _____. spontaneous fission; Easy 90. _____ is the amount of time it takes for one-half of a radioactive isotope to decay.

Half-life; Easy 91. 1 decay/s = _____ 1 becquerel; Easy 92. _____ = 3.7  1010 decays/s 1 Ci; Easy 93. ________________ discovered natural radioactivity from uranium ores.

Henri Becquerel; Easy 94. _____= 100 rad

1 Gy; Easy 95. rem = _____  factor rad; Easy 96. A(n) _____has a tube with a thin window that allows radioactive emissions through.

Geiger counter; Easy 97. A(n)_____ is a substance that can be used to follow the pathway of that substance through a process. tracer; Easy 98. A(n) _____ is an apparatus designed to carefully control the progress of a nuclear reaction and extract the resulting energy for useful purposes. nuclear reactor; Easy 99. A single neutron can begin a fission process that grows exponentially in a phenomenon called a(n) _____. chain reaction; Easy 100. _____ is a nuclear process in which small nuclei are combined into larger nuclei, releasing energy.

Nuclear fusion; Easy

Chapter 16 Organic Chemistry True/ False 1. The simplest organic compounds are those composed of only two elements: carbon and oxygen.

False; Easy 2. Aliphatic hydrocarbons are hydrocarbons based on chains of C atoms.

True; Easy

3. Alkanes are aliphatic hydrocarbons with only single covalent bonds.

True; Easy 4. Alkenes are hydrocarbons that contain at least one C–C double bond.

True; Easy 5. Alkynes are hydrocarbons that contain at least a C–C triple bond.

True; Easy 6. Aromatic hydrocarbons have a special six-carbon ring called a benzene ring.

True; Easy 7. Ca(OH)2 is an organic compound.

False; Easy 8. Methane is the smallest alkane.

True; Moderate 9. A polymer is an example of a micromolecule, which is a name given to a large molecule.

False: Moderate 10. The compound below is not an aromatic compound.

False; Moderate 11. The correct name for the compound below is 4-ethylhexane.

C C C

C

C

C

C C

False; Moderate 12. CH3CH2CH2CH3 and CH3CH2CH (CH3)2 are isomers.

False; Moderate 13. 3,3,5-trimethylhexane has 6 carbon atoms. (Hint: Draw the structure.)

False; Moderate 14. The correct name for the compound below is 4,8-diethyl-2-nonene.

C C C

C

C

C

C

C

C

C

C C C

False; Moderate 15. The correct name for the compound below is 3-bromo-2-ethylbenzene.

H H

C

Br

C

C

C

C C

H

H H

C

H

C

H H

H

False; Moderate 16. The correct name for the compound below is 2,3-diiodohexane.

C

I

I

C

C

C

C

C

True; Moderate 17. The correct name for the compound below is 6-heptanone.

O C

C

C

C

False; Moderate

C

C

C

18. The correct name for the molecule below is butyl ethanoate.

O C

C

C

C

False; Moderate

O

C

C

19. The functional group present in the molecule below is an ether.

O C

C

C

C

False; Moderate

C

OH

20. The functional group present in the molecule below is an ester.

C C C C C O C C False; Moderate 21. The bond between the N of the amine group and the C of the carbonyl group is called an amide bond.

True; Moderate 22. An amine is an organic derivative of sulfur.

False; Moderate 23. R–SH is the general formula of an aldehyde.

False; Moderate 24. A long, almost nonstop molecule is called a polymer.

True; Moderate 25. The sulfur analog of an alcohol is called a peptide.

False; Moderate

Multiple Choice Questions 26. The alkane with the formula, C2H6, is _____. ee. ethane ff. propane gg. methane hh. butane ii. pentane a; Easy 27. CH3(CH2)7CH3 is the condensed structural formula of _____. a. ethane b. propane c. heptane d. octane e. nonane e; Easy

28. The condensed structural formula of hexane is _____. a. CH3(CH2)3CH3 b. CH3(CH2)4CH3 c. CH3(CH2)5CH3 d. CH3(CH2)6CH3 e. CH3(CH2)7CH3 b; Easy 29. Which of the following is the smallest alkene? a. methene b. ethene c. propene d. butene e. pentene b; Easy 30. What is the proper name for the molecule shown below?

H

H

H

H

C

C

C

C

H

H

H

H

H

H

C

C

C

H

H

H

a. 4-heptene b. 3,4-heptene c. 3-heptane d. 4-heptane e. 3-heptene e; Easy 31. The structural formula of 1-hexene is _____. a.

H

H

H

H

H

H

C

C

C

C

C

H

H

H

H

H

H

H

H

C

C

C

C

C

H

H

H

H

H

H

H

H

H

C

C

C

C

C

C

H

H

H

H

H

H

H

b.

H

H

c.

H

d.

H

H

H

H

H

H

H

H

H

C

C

C

C

C

C

H

H

H

H

H

H

H

H

C

C

C

C

C

C

H

H

H

H

H

H

H

e.

H

H

c; Moderate 32. The smallest alkyne is _____. a. methyne b. butyne c. propyne d. ethyne e. pentyne d; Easy H H

C C

H

C

C

C C

H

H H

33.

is the structure of _____.

a. 2-hexene b. 2,3-hexene c. hexyne d. 3-hexene e. benzene e; Easy 34. What is the proper name for the molecule shown below?

H

H

H

H

C

C

C

H

H

H

a. 1-pentyne b. 5-pentyne c. 4-pentyne d. 4,5-pentane e. 4,5-pentene a; Moderate

C

C

H

35. Complete the reaction below.

H2 C

H3C

a.

H2 C

H3C

b.

H2C

H4 C

CH2

C H

CH2

H2 C

CH3

H2 C

H2 C

CH3

H2 C

H2 C

CH3

H2 C

H2 C

CH4

C H

c.

C H

+

H2

metal catalyst

d.

H3 C

e.

H3 C

d; Moderate 36. Complete the reaction below. light

? + Cl2  CH3Cl + HCl a. CH4 b. CH3 c. C2H6 d. C2H4 e. C3H5 a; Easy 37. Give the proper name for the molecule below.

C

C

C

C

C

C

C

a. 1,2-dimethylpentane b. 4,4-dimethylpentane c. 2,2-dimethylpentane d. 3,4-dimethylpentane e. 2,1-dimethylpentane c; Easy 38. Which of the following is the structure of 4-ethyl-3-methylhexane? a.

?

C C C

C

C

C

C

C

C

C

C

C

b.

C C C

C

C C C

c.

C C

C

C C C

d.

C C C

C

C

C

C C

e.

C C C

C C C

b; Moderate

C

C

39. Give the proper name for the molecule below.

C C C

C

C

C

C

C

C

C

C C C

a. 4,8-dipropyldecane b. 8-methyl-3-ethyldecane c. 3-propyl-7-ethylnonane d. 2-ethyl-7-propylnonane e. 3,7-diethyldecane e; Moderate 40. Give the proper name for the molecule below.

C C

C

C

C

C

C

C

C C C

a. 3,7-di-ethyl-5-nonene b. 2,7-diethyl-5-nonene c. 2,6-dimethyl-3-nonene d. 6-methyl-2-ethyl-3-nonene e. 2-propyl-7-methyl-6-nonene c; Moderate 41. Give the proper name for the molecule given below.

H H

C C

I C

C

C C

I H

H

C

a. 1,4-diiodobenzene b. 1,3-diiodobenzene c. 2,3-diiodobenzene d. 1,2-diiodobenzene e. 3,5-diiodobenzene a; Moderate 42. Give the proper name for the molecule below.

H F

C

H

C

C

C

C C

H

H H

C

H

H

C

H

H

C

H

H

a. 3-fluoro-1-propylbenzene b. 2-fluoro-3-propylbenzene c. 1-fluoro-3-propylbenzene d. 4-fluoro-1-propylbenzene e. 1-fluoro-5-propylbenzene c; Moderate 43. Give the proper name for the molecule below.

C C

C

C

C C

C

C

C

C

a. 6-pentylbenzene b. 1-pentylbenzene c. 4-phenylpentane d. 2-phenylpentane e. 5-phenylpentane d; Moderate

C

44. Give the proper name for the molecule below.

C

C

Cl

Br

C

C

C

a. 2-bromo-1-chloropentane b. 4-bromo-3-chloropentane c. 3-bromo-2-chloropentane d. 1-bromo-2-chloropentane e. 2-bromo-3-chloropentane e; Moderate 45. Give the proper name for the molecule below.

Cl C

C

Cl

Br

C

C

Br C

C

a. 3,1-dibromo-6,4-dichlorohexane b. 1,3-dibromo-4,6-dichlorohexane c. 2,3-dibromo-5,6-dichlorohexane d. 4,3-dibromo-4,5-dichlorohexane e. 4,3-dibromo-5,6-dichlorohexane b; Moderate 46. Give the proper name for the molecule below.

OH C

C

C

a. 5-propanol b. 4-propanol c. 2-propanol d. 3-propanol e. 1-propanol d; Easy

C

C

47. Give the proper name for the molecule below as a substituted alkane.

OH C

C

C

F C

C

C

Br

a. 4-bromo-6-fluoro-3-hydroxyhexane b. 4-hydroxy-3-bromo-1-fluorohexane c. 2-hydroxy-3-bromo-1-fluorohexane d. 4-hydroxy-5-bromo-6-fluorohexane e. 1-hydroxy-2-bromo-3-fluorohexane a; Moderate 48. Complete the reaction below.

H

H

H

H

OH

C

C

C

C

C

H

H

H

H

H

H

H

H

C

C

C

C

C

H

H

H

H

H

H

H

H

C

C

C

C

C

H

H

H

H

H

H

H

H

H

H

C

C

C

C

C

H

H

H

H

H

H

H

H

H

H

C

C

C

C

C

H

H

H

H

H

H

H

H

H

H

C

C

C

C

C

H

H

H

H

H

acid H

H

a.

H

H

b.

H

H

c.

H

H

d.

H

H

e.

H

H

b; Moderate 49. Give the proper name for the molecule below.

O C

C

C

a. 4-pentanone b. 2-pentanone c. 1-pentanone d. 3-pentanone e. 5-pentanone b; Easy

C

C

?

50. Give the proper name for the molecule given below.

O C

C

C

a. Butanal b. Butanone c. Butanoic acid d. Butanol e. Butyne a; Easy

C

H

51. Give the proper name for the molecule given below.

O C

C

C

C

a. Butyl methyl ketone b. Diethyl ketone c. Methyl propyl ketone d. Ethyl propyl ketone e. Dipropyl ketone e; Easy

C

C

C

52. The common name of methanal is _____. a. acetaldehyde b. formaldehyde c. acetic acid d. citric acid e. benzaldehyde b; Easy 53. Give the proper name for the molecule below.

O C

C

C

a. Pentanol b. Pentanal c. Pentanoate d. Pentanoic acid e. Pentanediol d; Easy

C

C

OH

54. Complete the chemical reaction.

O C

C

C

C

OH

a.

O C

C

C

C

H-

+

OH-

?

+

H2O

b.

H C

C

C

C

OH

c.

O C

d.

C

C

C

OH

H C

e.

C

C

C

OOH+

O C

C

e; Moderate

C

C

O-

55. Complete the chemical reaction.

O C

C

C

OH

+

C

C

OH

a.

O C

b.

C

C

OO-

O C

C

C

OOH

c.

O C

C

C

C

O

C

C

C

C

C

C

O

C

C

d. e.

O C

C

C

C

H

c; Moderate 56. Give the proper name for the molecule below.

O C

C

C

C

a. Ethyl pentyl ketone b. Pentyl ethanal

C

O

C

C

?

+

H2O

c. Ethyl pentyl ether d. Ethyl pentanoate e. Methyl butanoic acid

d; Moderate 57. Give the proper name for the molecule below.

C C C C C O C C a. Ethyl pentyl ketone b. Pentyl ethanal c. Ethyl pentyl ether d. Ethyl pentanoate e. Methyl butanoic acid

c; Moderate 58. A functional group that has an O atom attached to two organic groups. a. amine b. ether c. ester d. carboxylic acid e. ketone b; Easy 59. Give the proper name for the molecule below.

C C C

C

N

C

C

a. Methyl propyl amine b. Ethyl methyl propyl amine c. Ethyl methyl amine d. Methyl ethyl amine e. Triethyl amine e; Easy 60. Give the proper name for the molecule below.

H N

H

a. Phenal b. Phenol c. Benzylamide

d. Phenylamine e. Benzylamine d; Moderate 61. The bond between the N of the amine group and the C of the carbonyl group is called a(n)_____ bond. a. phenol b. alcohol c. aldehyde d. amide e. ketone d; Easy 62. Give the proper name for the molecule below.

C

C

O

H

C

N

a. Phenal propanamide b. Phenyl propanamide c. Propyl benzylamide d. Benzyl butanamide e. Benzyl proanamide b; Moderate 63. The sulfur analog of an alcohol is called a(n) _____. a. aldehyde b. ketone c. thiol d. ester e. ether c; Easy 64. Which term was used in the past to denote simple sulfur-containing hydrocarbons? a. aldehyde b. ketone c. mercaptan d. ester e. ether c; Easy 65. Give the proper name for the molecule below.

C C C C C SH a. Pentanethiol b. Pentanediol c. Pentanoate d. Pentanone e. Pentanal a; Easy

66. Draw the polymer made from this monomer.

H

H

C

C

I

H

a.

H

I

H

H

C

C

C

C

H

H

H

H

I

I

I

I

C

C

C

C

H

H

H

H

H

H

H

H

C

C

C

C

I

H

I

H

H

H

H

H

C

C

C

C

H

H

H

H

n

b.

n

c.

n

d.

n

c; Easy 67. _____ is a reaction in which small molecules, referred to in general as monomers, are assembled into giant molecules. a. Alkylation b. Halogenation c. Hydrogenation d. Polymerization e. Combustion d; Easy

H Si

H O

Si

O

n

68.

is a(n) _____. a. salt b. silicone c. thiol d. mercaptan e. diol b; Easy

69. ______ is used in the manufacture of polyethylene, one of the most familiar plastics. a. Methylene b. Ethylene c. Butyne d. Pentyne e. Propylene b; Easy 70. The condensed structural formula for the section of a molecule formed from five units of the monomer CHI=CH2 is: a. ~CH2ICH2CH3ICH2CH2I2CH2CHI2CH2CH3ICH3~. b. ~CH2ICH2CH2ICH2CH2ICH2CH2ICH2CH2ICH2~. c. ~CHI2CH2CHI2CH2CHICH2CHI3CH2CHI3CH2~. d. ~CHICH2CHICH2CHICH2CHICH2CHICH2~. e. ~CHICH2CHI3CH2CHI2CH2CH2ICH2CHICH2~. d; Moderate Essay Questions 71. Differentiate between aliphatic and aromatic hydrocarbons. What makes alkanes, alkenes, alkynes, different?

Aliphatic hydrocarbons are hydrocarbons based on chains of C atoms. Aromatic hydrocarbons have a special six-carbon ring called a benzene ring. Originally, the term aromatic was used to describe this class of compounds because they were particularly fragrant. However, in modern chemistry the term aromatic denotes the presence of a sixmembered ring that imparts different and unique properties to a molecule.

There are three types of aliphatic hydrocarbons. Alkanes are aliphatic hydrocarbons with only single covalent bonds. Alkenes are hydrocarbons that contain at least one C–C double bond, and alkynes are hydrocarbons that contain a C–C triple bond. Occasionally, we find an aliphatic hydrocarbon with a ring of C atoms; these hydrocarbons are called cycloalkanes (or cycloalkenes or cycloalkynes).

Moderate 72. What are isomers?

Different molecules with the same molecular formula are called isomers. Isomers are common in organic chemistry and contribute to its complexity.

Easy 73. List some properties of hydrocarbons.

Most hydrocarbons are nonpolar because of the close electronegativities of the C and H atoms. As such, they dissolve only sparingly in H2O and other polar solvents.

Halogens can also react with alkenes and alkynes, but the reaction is different. In these cases, the halogen reacts with the C–C double or triple bond and inserts itself onto each C atom involved in the multiple bonds. This reaction is called an addition reaction. One example is

Hydrogen can also be added across a multiple bond; this reaction is called a hydrogenation reaction.

CH 2  CH 2  H 2

metal catalyst



CH3CH3

By far the most common reaction of hydrocarbons is combustion, which is the combination of a hydrocarbon with O2 to make CO2 and H2O. 2C8H18 + 25O2 → 16CO2 + 18H2O + ~5060 kJ Moderate 74. Describe the IUPAC nomenclature.

There are a series of rules for naming branched alkanes (and, ultimately, for all organic compounds). These rules make up the system of nomenclature for naming organic molecules.

Consider this molecule:

The longest chain has four C atoms, so it is butane. There are two substituents, each of which consists of a single C atom; they are methyl groups. The methyl groups are on the second and third C atoms in the chain (no matter which end the numbering starts from), so we would name this molecule 2,3-dimethylbutane. Note the comma between the numbers, the hyphen between the numbers and the substituent name, and the presence of the prefix di- before the methyl. Other molecules—even with larger numbers of substituents—can be named similarly.

Moderate 75. Define a functional group.

A functional group is any collection of atoms and/or bonds with certain characteristic chemical reactions. We have already seen two functional groups: the C–C double bond and the C–C triple bond. They undergo certain characteristic chemical reactions—for example, the addition of a halogen across the multiple bond.

Moderate

76. Identify and name a simple alkyl halide.

The presence of a halogen atom (F, Cl, Br, or I; also, X is used to represent any halogen atom) is one of the simplest functional groups. Organic compounds that contain a halogen atom are called alkyl halides.

A simple alkyl halide can be named like an ionic salt, first by stating the name of the parent alkane as a substituent group (with the -yl suffix) and then the name of the halogen as if it were the anion. So CH3Cl has the common name of methyl chloride, while CH3CH2Br is ethyl bromide and CH3CH2CH2I is propyl iodide. However, this system is not ideal for more complicated alkyl halides.

Moderate 77. Identify and name a simple alcohol.

An alcohol is a simple functional group that is the covalently bonded OH group. This is the alcohol functional group. It is not the hydroxide ion; rather than being present as a negatively charged species, in organic chemistry it is a covalently bonded functional group.

Alcohols have a common naming system and a formal system. The common system is similar to that of alkyl halides: name the alkyl group attached to the OH group, ending with the suffix -yl, and add the word alcohol as a second word. So CH3OH is methyl alcohol; CH3CH2OH is ethyl alcohol, and CH3CH2CH2OH is propyl alcohol.

Moderate 78. Predict the product(s) of an elimination reaction of an alkyl halide or an alcohol.

One reaction common to alcohols and alkyl halides is elimination, the removal of the functional group (either X or OH) and an H atom from an adjacent carbon. The general reaction can be written as follows:

where Z represents either the X or the OH group. The biggest difference between elimination in alkyl halides and elimination in alcohols is the identity of the catalyst: for alkyl halides, the catalyst is a strong base; for alcohols, the catalyst is a strong acid. For compounds in which there are H atoms on more than one adjacent carbon, a mixture of products results.

Moderate 79. Identify the aldehyde and ketone.

First the carbonyl group is defined. A crabonyl group is formed when an O atom and a C atom are joined by a double bond.

O C

If one bond of the carbonyl group is made to an H atom, then the molecule is classified as an aldehyde. (If there are two H atoms, there is only 1 C atom.)

When naming aldehydes, the main chain of C atoms must include the carbon in the carbonyl group, which is numbered as position 1 in the carbon chain. The parent name of the hydrocarbon is used, but the suffix -al is appended.

A carbonyl group in the middle of a carbon chain implies that both remaining bonds of the carbonyl group are made to C atoms. This type of molecule is called a ketone.

When naming a ketone, we take the name of the parent hydrocarbon and change the suffix to –one.

Moderate 80. Identify the carboxylic acid and ether.

Molecules with a carboxyl group are called carboxylic acids. As with aldehydes, the functional group in carboxylic acids is at the end of a carbon chain. Also as with aldehydes, the C atom in the functional group is counted as one of the C atoms that defines the parent hydrocarbon name. To name carboxylic acids, the parent name of the hydrocarbon is used, but the suffix -oic acid is added. the ether functional group is an O atom that is bonded to two organic groups: R—O—R′ The two R groups may be the same or different. Naming ethers is like the alternate way of naming ketones. In this case, the R groups are named sequentially, and the word ether is appended.

Moderate 81. Identify an amine.

An amine is an organic derivative of ammonia (NH3). In amines, one or more of the H atoms in NH3 is substituted with an organic group. Naming simple amines is straightforward: name the R groups as substituents and then add the suffix -amine, using numerical suffixes on the substituent names as necessary.

Moderate 82. Describe the classification of amines.

A primary amine has one H atom substituted with an R group:

A secondary amine has two H atoms substituted with R groups:

A tertiary amine has all three H atoms substituted with R group:

Moderate 83. Identify an amide.

An amide functional group is a combination of an amine group and a carbonyl group.

Amides are actually formed by bringing together an amine-containing molecule and a carboxylic acid-containing molecule.

The bond between the N of the amine group and the C of the carbonyl group is called an amide bond.

Moderate 84. Identify a thiol.

The sulfur analog of an alcohol is called a thiol. The formal way of naming a thiol is similar to that of alcohols, except that instead of using the suffix -ol, you use -thiol as the suffix.

Moderate 85. Describe polymerization.

Consider a molecule with a double bond, such as ethylene: CH2=CH2 Imagine the bond between the carbons opening up and attacking another ethylene molecule.

–CH2–CH2– Now imagine further that the second ethylene molecule‘s double bond opens up and attacks a third ethylene molecule, which also opens up its double bond and attacks a fourth ethylene molecule, and so forth. The end result is long, virtually endless molecule: –CH2–CH2–CH2–CH2– This long, almost nonstop molecule is called a polymer. The original part—ethylene—is called the monomer (meaning ―one part‖). The process of making a polymer is called polymerization. A polymer is an example of a macromolecule, the name given to a large molecule.

Moderate Fill in the Blanks 86. An organic compound composed of carbon and hydrogen is called a(n)_____. hydrocarbon; Easy 87. Most hydrocarbons are __________ because of the close electronegativity of C and H atoms. nonpolar; Easy

88. A molecule with the same molecular formula as another molecule but a different structure is called a(n)_____. isomer; Easy 89. _________________ is a compound with less than the maximum possible number of H atoms in its formula.

Unsaturated hydrocarbon; Easy 90. The reaction of a hydrogen molecule across a C–C double or triple bond is called a(n)_____.

hydrogenation reaction; Easy 91. A _____ is formed when an O atom and a C atom are joined with a double bond.

carbonyl group; Easy 92. A(n) _____ is a negatively charged ion derived from a carboxylic acid. carboxylate ion; Easy 93. _____ is a functional group made by combining a carboxylic acid with an alcohol.

Ester; Easy 94. A(n) _____ is a functional group that has an O atom attached to two organic groups.

Ether; Easy 95. A(n) _____ is a compound that has a carbonyl group in the middle of a carbon chain.

Ketone; Easy 96. A(n) _____ is an organic derivative of ammonia.

Amine; Easy 97. The bond between the N atom and the C atom in an amide is called a(n)_____. amide bond; Easy 98. Most common reaction of hydrocarbons is _____________. combustion; Easy 99. A(n)_____ is a long molecule made of hundreds or thousands of repeating units. polymer; Easy 100. The repeated unit of a polymer is a(n)_____. monomer; Easy