hhs eoc chemistry test prep -...

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HHS

EOC Chemistry

Test Prep

Name: ____________________________________

Note: These are just practice questions…studying and

reviewing your notes will be beneficial in helping prepare you

for the EOC Chemistry Exam.

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Measurement

Unit Conversions- Kilo, Hecto, Deka, Base unit, Deci, Centi, Milli.

Steps of the Scientific Method

Make observation- using your senses

State the problem

Form a hypothesis

Design/ perform an experiment

Collect/ analyze the data

Make predictions based on the results

Make adjustments and repeat the experiment

Key Terms

Control- factor in an experiment that is purposely kept the same

Data table- evidence or information gathered from observations in an organized format

Graph- a drawing or diagram that shows a relationship between sets of things

o Independent variable- on the x-axis

o Dependent variable- on the y-axis

Hypothesis- possible explanation for a set of observations or possible answer to a

scientific question

Mean- an average

Median- the middle number

Proportionality- the relationship of a part to a whole

Variable- anything that can change in an experiment

Percent Composition- Percent is ‘part divided by whole’. To calculate percent composition, find

the molar mass for each atom with the compound. Take each individual molar mass (of the

element) and divide by the total molar mass of the compound.

1. 0.25g is equivalent to

a. 250kg

b. 250mg

c. 0.025mg

d. 0.025kg

2. 0.05cm is the same as

a. 0

b. .000 05m

c. 0.005mm

d. 0.05m

e. 0.5mm

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3. 1.06L of water is equivalent to

a. 0.001 06mL

b. 10.6mL

c. 106mL

d. 1060mL

4. The number of grams equal to 0.5kg is

a. 0.0005 g

b. 0.005 g

c. 500 g

d. 5000 g

5. What is the percentage composition of CF4?

a. 20% C, 80% F

b. 13.6% C, 86.4% F

c. 16.8% C, 83.2% F

d. 81% C, 19% F

6. What is the percentage composition of CO?

a. 50% C, 50% O

b. 12% C, 88% O

c. 25% C, 75% O

d. 43% C, 57% O

7. What is the percentage composition of CuCl2?

a. 33% Cu, 66% Cl

b. 50% Cu, 50% Cl

c. 65.50% Cu, 34.50% Cl

d. 47.27% Cu, 52.73% Cl

8. The percentage composition of sulfur in SO2 is about 50%. What is the percentage of

oxygen in this compound?

a. 25%

b. 50%

c. 75%

d. 90%

9. All of the following are steps in the scientific method except

a. observing and recording data.

b. forming a hypothesis.

c. discarding data inconsistent with the hypothesis.

d. developing a model based on experimental results.

10. The initial or tentative explanation of an observation is called a(n)

a. law.

b. theory.

c. hypothesis.

d. experiment.

11. Which of the following observations is quantitative?

a. The liquid turns blue litmus

paper red.

b. The liquid boils at 100°C.

c. The liquid tastes bitter.

d. The liquid is cloudy.

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12. Which of the following observations is qualitative?

a. A chemical reaction was complete in 2.3 seconds.

b. The solid had a mass of 23.4 g.

c. The pH of a liquid was 5.

d. Salt crystals formed as the liquid evaporated.

13. Which of the following has the same number of significant figures as the number

1.00310?

a. 1 × 106

b. 199.791

c. 8.66

d. 5.119

14. The number with the most significant zeroes is

a. 0.00002510

b. 0.02500001

c. 250000001 d. 2.501 × 10-7

15. How many significant figures should be retained in the result of the following

calculation?

12.00000 × 0.9893 + 13.00335 × 0.0107

a. 2

b. 3

c. 4

d. 5

16. The correct answer (reported to the proper number of sig. figs.) to the following is

6.3 × 3.25

a. 20.

b. 20.475

c. 20.48

d. 20.5


(2.05631) (6.9391136) / 12.59326

a. 1.1330639

b. 1.13306

c. 1.133064

d. 1.1361


12.75 × 1.3621

a. 17.367

b. 17.40

c. 17.37

d. 17.4


(12.67 + 19.2)(3.99) / (1.36 + 11.366)

a. 9.99851

b. 9.9985

c. 9.999

d. 1.00 × 101

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20. How is the measurement 0.000065cm written in scientific notation?

a. 65 × 10-6 cm

b. 6.5 × 10-5 cm

c. 6.5 × 10-6 cm

d. 6.5 × 10-4 cm

21. How would 0.00930m be expressed in scientific notation?

a. 93 × 10-4 m

b. 9.3 × 10-4 m

c. 9.30 × 10-3 m

d. 9.30 × 10-5 m

22. An analytical balance can measure mass to the nearest 0.0001g. In scientific notation,

the accuracy of the balance would be expressed as

a. 1.0 × 10-3 g.

b. 1 × 103 g.

c. 1 × 104 g.

d. 1 × 10-4 g.

23. When 1.92 × 10-6 kg is divided by 6.8 × 102 mL, the quotient equals

a. 2.8 × 10-4 kg/mL.

b. 2.8 × 10-5 kg/mL.

c. 2.8 × 10-8 kg/mL.

d. 2.8 × 10-9 kg/mL.

24. When 6.02 × 1023 is multiplied by 9.1 × 10-31, the product is

a. 4.3 × 10-8

b. 4.3 × 1054

c. 4.3 × 10-7

d. 4.3 × 10-53

25. The results of dividing 107 by 10-3 is

a. 10-4

b. 102.5

c. 104

d. 1010

26. Choose the number that is most accurate.

a. 5.0 m

b. 5 m

c. 5.25 m

d. 5.250 m

Calculations

Molarity- a concentration unit of a solution expressed as moles of solute dissolved per liter of

solution.

Molality- the concentration of a solution expressed in moles of solute per kilogram of solvent

27. What is the molarity of a solution that contains 0.202 mol KCl in a 7.98 solution?

a. 0.0132 M

b. 0.0253 M

c. 0.459 M

d. 1.363 M

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28. What is the molality of a solution that contains 5.10 mol KNO3 in 4.47kg water? (molar

mass of KNO3 = 101.11 g/mol)

a. 0.315 m

b. 0.779 m

c. 1.02 m

d. 1.14 m

29. What is the molarity of a solution that contains 125g NaCl in 4.00L solution? (molar mass

of NaCl = 58.44 g/mol)

a. 0.535 M

b. 2.14 M

c. 8.56 M

d. 31.3 M

30. How many oxygen atoms are there in 0.500 mol of CO2?

a. 6.02 × 1023

b. 3.01 × 1023

c. 15.9994

d. 11.0

31. How many Mg2+ ions are found in 1.00 mol of MgO?

a. 3.01 × 1023

b. 6.02 × 1023

c. 12.04 × 1023

d. 6.02 × 1025

32. If 0.500 mol of Na+ combines with 0.500 mol of Cl- to form NaCl, how many atoms of

NaCl are present?

a. 3.01 × 1023

b. 6.02 × 1023

c. 6.02 × 1024

d. 1

33. How many molecules are there in 5.0g of methyl alcohol, CH3OH?

a. 9.4 × 1022

b. 3.0 × 1024

c. 3.6 × 1024

d. 3.8 × 1024

Density, pH

Density- the ratio of the mass of a substance to the volume of the substance. It is often

expressed as grams per cubic centimeter for solids and liquids and as grams per liter for gases.

Mass- a measure of the amount of matter in an object. It does not change.

Weight- a measure of the gravitational force exerted on an object. Its value can change with

the location of the object in the universe.

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pH- a value that is used to express the acidity or alkalinity (basicity) of a system. Each whole

number on the scale indicated a tenfold change in acidity. A pH of 7 is neutral, a pH of less than

7 is acidic, and a pH of greater than 7 is basic.

pH Measurement Scale

Neutral

increasing acidity <- -> increasing basic

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

* * * * *

Stomach Tomatoes H2O Soap Lye

Acid

34. Which pair of quantities determines the density of a material?

a. Mass and weight

b. Volume and weight

c. Volume and concentration

d. Volume and mass

35. The density of an object is calculated by

a. multiplying its mass times its

volume.

b. dividing its mass by its volume.

c. dividing its volume by its mass.

d. adding its mass to its volume.

36. The density of aluminum is 2.70 g/cm3. What is the mass of a solid piece of aluminum

with a volume of 1.50 cm3?

a. 0.556 g

b. 1.80 g

c. 4.05 g

d. 4.20g

37. A sample of gold has a mass of 96.5g and a volume of 5.00 cm3. The density of gold is

a. 0.0518 g/cm3

b. 19.3 g/cm3

c. 101.5 g/cm3

d. 483 g/cm3

38. The density of pure diamond is 3.5 g/cm3. What is the volume of a diamond with a mass

of 0.25g?

a. 0.071 cm3

b. 0.875 cm3

c. 3.75 cm3 d. 14 cm3

39. What is the density of 37.72g of material whose volume is 6.80 cm3?

a. 0.180 g/cm3

b. 5.55 g/cm3

c. 30.9 g/cm3

d. 256 g/cm3

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40. What is the pH of a neutral solution at 25°C?

a. 0

b. 1

c. 7

d. 14

41. The pH scale, in general, uses ranges from

a. 0 to 1.

b. -1 to 1.

c. 0 to 7.

d. 0 to 14.

42. If [H3O+] = 8.26 × 10-5 M, what is the pH of the solution?

a. 2.161

b. 3.912

c. 4.083

d. 8.024

43. What is the pH of a solution whose hydronium ion concentration is 5.03 × 10-1 M?

a. 0.2984

b. 0.5133

c. 1.542

d. 5.031

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Mixtures

A mixture is two or more substances together which can be separated by physical means.

If the mixture has a uniform composition, it is said to be homogeneous- like salt water.

If the composition of the mixture is not uniform, it is said to be heterogeneous- like Italian

salad dressing.

44. A combination of sand, salt and water is an example of a

a. homogeneous mixture.

b. heterogeneous mixture.

c. compound.

d. pure substance.

Energy

45. Differentiate between potential (stored) and kinetic (moving) energy.

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Acids and Bases

46. List the most common acids.

47. List the most common bases.

48. What is an Arrhenius acid?

49. What is a Bronsted-Lowry acid?

50. What is an example of a Bronsted-Lowry base?

Periodic Table

51. Because most particles fired at metal foil passed straight through, Rutherford

concluded that

a. atoms were mostly empty space.

b. atoms contained no charge

particles.

c. electrons formed the nucleus.

d. atoms were indivisible.

52. Rutherford’s experiments led him to conclude that atoms contain massive central regions

that have

a. a positive charge.

b. a negative charge.

c. no charge.

d. both protons and electrons.

53. The nucleus of an atom has all of the following characteristics except that it

a. is positively charged.

b. is very dense.

c. contains nearly all of the atom’s mass.

d. contains nearly all of the atom’s volume.

54. An atom is electrically neutral because

a. neutrons balance the protons and electrons.

b. nuclear forces stabilize the charges.

c. the number of protons and electrons are equal.

d. the numbers of protons and neutrons are equal.

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55. Most of the volume of an atom is occupied by the

a. nucleus.

b. nuclides.

c. electrons.

d. protons.

56. The charge due to the electrons of a neutral atom

a. prevents compounds from

forming.

b. balances the charge on the

nucleus.

c. attracts electrons in other

atoms.

d. does not exist.

57. What are the radioactive elements with atomic numbers from 90 to 103 called?

a. The noble gases

b. The lanthanides

c. The actinides

d. The rare-earth elements

58. What are the elements with atomic numbers from 58 to 71 called?

a. The lanthanides

b. The noble gases

c. The actinides

d. The alkali metals

59. Argon, krypton and xenon are

a. alkaline earth metals.

b. noble gases.

c. actinides.

d. lanthanides.

60. Elements in a group or column in the periodic table can be expected to have similar

a. atomic masses.

b. atomic numbers.

c. numbers of neutrons.

d. properties.

61. The most reactive group of the nonmetals is the

a. lanthanides.

b. transition elements.

c. halogens.

d. noble gases.

62. The group soft, silvery, reactive metals, all of which have one electron in an s orbital, is

known as the

a. alkaline-earth metals.

b. transition metals.

c. alkali metals.

d. metalloids.

63. The first member of the noble gas family, whose highest energy level consists of an

octet of electrons, is

a. Helium. b. Argon.

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c. Neon. d. Krypton.

64. The most characteristic property of the noble gases is that they

a. have low boiling points.

b. are radioactive.

c. are gases at ordinary temperatures.

d. are largely unreactive.

65. Nonmetals are

a. shiny.

b. poor electrical conductors.

c. malleable.

d. ductile.

66. Molecules have melting points that are

a. Close to melting points of ionic compounds.

b. Lower than the melting points of ionic compounds.

c. Higher than the melting points of ionic compounds.

d. Based on the polarity of the molecule.

Periodic Trends

Atomic radius- one-half of the distance between the center of identical atoms that are not

bonded together.

Ionization energy- the energy required to remove an electron from an atom or ion. The smaller

the atom, the closer the valence electrons are to the nucleus and therefore, the tighter the

electrons are being held. This gives the smallest atoms the largest ionization energy. It should

be intuitive that as you go down a family/group (column), the atomic radius gets larger because

you add in another energy level for the electrons for each space you go down. This means as you

go down the family, the ionization energy will get smaller. Think about a football player running

with the ball. If he holds it in tight against his body, it is difficult to get it away from him.

However, if he is running down the field with his arm outstretched holding the ball it is easy to

take the ball away from him (and his coach is having a heart attack). What may not be intuitive

is that the atomic radius actually gets smaller as you go across a period (row). You would think

that if you add in electrons the radius should get bigger, but it doesn’t. You are also adding in a

proton. As you add in an electron, which adds mass and increases the nuclear force of the

nucleus, so it can hold the electrons even tighter. So the largest ionization energies are in the

upper right of the periodic table and the lowest are in the bottom left.

67. Within a group of elements, as the atomic number increases, the atomic radius

a. increases.

b. remains approximately constant.

c. decreases regularly.

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d. varies unpredictably.

68. In the alkaline-earth metals, atoms with the smallest radii

a. are the most reactive.

b. have the largest volume.

c. are all gases.

d. have the highest ionization energies.

69. Across a period in the periodic table, atomic radii

a. gradually decrease.

b. gradually decrease, then sharply increase.

c. gradually increase.

d. gradually increase, then sharply decrease.

70. Which is the best reason that the atomic radius generally increases with atomic number

in each group of elements?

a. The nuclear charge increases.

b. The number of neutrons increases.

c. The number of occupied energy levels increases.

d. A new octet forms.

71. The energy required to remove an electron from an atom is the atom’s

a. electron affinity.

b. electron energy.

c. electronegativity.

d. ionization energy.

72. A measure of the ability of an atom in a chemical compound to attract electrons from

another atom in the compound is called

a. electron affinity.

b. electron configuration.

c. electronegativity.

d. ionization potential.

73. One-half the distance between the nuclei of identical atoms that are bonded together is

called the

a. atomic radius.

b. atomic diameter.

c. atomic volume.

d. electron cloud.

74. The element that has the greatest electronegativity is

a. Oxygen.

b. Sodium.

c. Chlorine.

d. Fluorine.

75. Electronegativity values

a. increase from left to right.

b. decrease from left to right.

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c. increase from top to bottom.

d. are calculated based on isotope number.

76. Which molecule can conduct electricity when placed in water?

a. C2H4

b. C6H12O6

c. NaNO3

d. H2

77. Which compound is soluble in water?

a. NaNO3

b. CF4

c. BaSO4

d. CaCO3

Isotopes

Isotopes- two atoms that have different numbers of neutrons but the same number of protons

and electrons. If they have the same number of protons but different number of neutrons, they

must have different masses.

78. The Cl- atom has an atomic number 17 and mass number 35. It has

a. 17 protons, 16 electrons, and 18 neutrons.

b. 35 protons, 35 electrons, and 17 neutrons.

c. 17 protons, 17 electrons, and 52 neutrons.

d. 18 protons, 18 electrons, and 17 neutrons.

79. Carbon-14 (atomic number 6), the radioactive nuclide used in dating fossils, has

a. 6 neutrons.

b. 8 neutrons.

c. 10 neutrons.

d. 14 neutrons.

80. Phosphorus-33 (atomic number 15) contains

a. 33 protons.

b. 18 neutrons.

c. 33 neutrons.

d. 18 protons.

81. Neon-22 contains 12 neutrons. It also contains

a. 12 protons.

b. 22 protons.

c. 22 electrons.

d. 10 protons.

82. The total number of protons and neutrons in the nucleus of an atom is its

a. atomic number.

b. Avogadro number.

c. mass number.

d. average atomic mass.

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83. As the mass number of an element’s isotopes of an element increases, the number of

protons

a. decreases.

b. increases.

c. remains the same.

d. doubles each time the mass number increases.

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Bonding

Ionic bonds- happen between metals and non-metals. Ionic bonds happen between two elements

that have oppositely charged ions. Ionic bonding requires a transfer of electrons. Use the criss-

cross method to determine chemical compounds. Transition elements require the use of Roman

Numerals in naming compounds. Polyatomic ions are compounds that contain two or more

elements combined that act as one.

Covalent bonds- happen between non-metals and non-metals. Covalent bonds share electrons.

When naming covalent bonds, prefixes must be used.

84. A positive ion is known as a(n)

a. ionic radius.

b. valence electron.

c. cation- electrons are removed.

d. anion.- electrons are added.

85. A negative ion is known as a(n)

a. ionic radius.

b. valence electron.

c. cation.

d. anion.

86. What are shared in a covalent bond?

a. Ions

b. Lewis structures

c. Electrons

d. Dipoles

87. Most chemical bonds are

a. purely ionic.

b. purely covalent.

c. partly ionic and partly covalent.

d. metallic.

88. What is the formula for zinc (II) fluoride?

a. ZnF

b. ZnF2

c. Zn2F

d. Zn2F3

89. What is the formula for the compound by calcium ions and chloride ions?

a. CaCl

b. Ca2Cl

c. CaCl3

d. CaCl2

90. What is the formula for the compound formed by lead (II) ions and chromate ions?

a. PbCrO4

b. Pb2CrO4

c. Pb2(CrO4)3

d. Pb(CrO4)2

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91. What is the formula for aluminum sulfate?

a. AlSO4

b. Al2SO4

c. Al2(SO4)3

d. Al(SO4)3

92. What is the formula for barium hydroxide?

a. BaOH

b. BaOH2

c. Ba(OH)2

d. Ba(OH)

93. Name the compound Ni(ClO3)2.

a. Nickel (II) chlorate

b. Nickel (II) chloride

c. Nickel (II) chlorite

d. Nickel (II) peroxide

94. Name the compound Zn3(PO4)2.

a. Zinc potassium oxide

b. Trizinc polyoxide

c. Zinc (II) phosphate

d. Zinc phosphite

95. What is the correct Lewis structure for hydrogen chloride, HCl?

a. A

b. B

c. C

d. D

96. If the atoms that share electrons have an unequal attraction for the electrons, the bond

is called

a. nonpolar.

b. polar.

c. ionic.

d. dipolar.

97. Which bond is most polar?

a. H- F

b. Br- Br

c. H- Ca

d. O- F

98. The B-F bond in BF3 (electronegativity for B is 2.0 and F is 4.0) is

a. polar covalent.

b. ionic

c. nonpolar covalent.

d. metallic.

99. A neutral group of atoms held together by covalent bonds (non-metals) is a

a. molecular formula.

b. chemical formula.

c. polyatomic ion.

d. molecule.

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100. Which of the following shows the types and numbers of atoms joined in a single

molecule of a molecular compound?

a. Molecular formula

b. Potential energy diagram

c. Covalent bond

d. Ionic bond

101. Which of the following is not an example of a molecular formula?

a. H2O

b. B

c. NH3

d. O2

102. A chemical bond formed by the attraction between positive ions and surrounding

mobile electrons is a(n)

a. nonpolar covalent bond.

b. ionic bond.

c. polar covalent bond.

d. metallic bond.

103. VSEPR theory is a model for predicting

a. the strength of metallic bonds.

b. the shape of the molecules.

c. lattice energy values.

d. ionization energy.

104. The concept that electrostatic repulsion between electron pairs surrounding an

atom causes these pairs to be separated as far as possible is the foundation of

a. VSEPR theory.

b. the hybridization model.

c. the electron-sea model.

d. Lewis theory.

105. Which of the following groups of ions are written correctly?

a. Ca+2, B+3, Cl+1

b. H+, N+2, C+7

c. Li+, Si-, Se+3

d. Be+2, Br-, S-2

e- configuration

Valence electrons- the electrons on the outermost electron shell of each atom. The number of

valence electrons can never exceed 8 (octet rule).

106. In the electron configuration for scandium (atomic number 21), what is the

notation for the three highest-energy electrons?

a. 4s2 3d1

b. 4s3

c. 3d3

d. 4s2 4p1

107. The element with electron configuration 1s2 2s2 2p6 3s2 3p2 is

a. Mg (Z = 12).

b. C (Z = 6).

c. S (Z = 16).

d. Si (Z = 14).

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108. The electron configuration for the carbon atom is 1s2 2s2 2p2. The atomic number

of carbon is

a. 3

b. 6

c. 11

d. 12

109. What is the electron configuration for nitrogen, atomic number 7?

a. 1s2 2s2 2p3

b. 1s2 2s3 2p2

c. 1s2 2s3 2p1

d. 1s2 2s2 2p2 3s1

110. Which element has the following electron configuration: [Ar] 4s2 3d10 4p5?

111. Write the noble gas electron configuration for Silicon.

112. Draw the orbital diagram for Phosphorus.

113. Draw the orbital diagram for Argon.

114. Write the noble gas electron configuration represented in the orbital diagram

below.

Gases

115. The volume of a gas is 400.0mL when the pressure is 1.00 atm. At the same

temperature, what is the pressure at which the volume of the gas is 2.0 L?

a. 0.5 atm

b. 5.0 atm

c. 0.20 atm

d. 800 atm

116. A sample of oxygen occupies 560.0 mL when the pressure is 800.00 mmHg. At

constant temperature, what volume does the gas occupy when the pressure decreases to

700.0 mmHg?

a. 80.0 mL

b. 490 mL

c. 600 mL

d. 640 mL

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117. The volume of a sample of oxygen is 300.0mL when the pressure is 1.00 atm and

the temperature is 27.0°C. At what temperature is the volume 1.00 L and the pressure

0.500 atm?

a. 22.0°C

b. 45.0°C

c. 0.50°C

d. 227°C

118. The volume of a gas collected when the temperature is 11.0°C and the pressure is

710mmHg measures 14.8 mL. What is the calculated volume of the gas at 20.0°C and 740

mmHg?

a. 7.8 mL

b. 13.7 mL

c. 14.6 mL

d. 15 mL

119. What determine the average kinetic energy of the molecules of any gas?

a. Temperature

b. Pressure

c. Container volume

d. Molar mass

120. The molar volume of a gas at STP is ______ L.

a. 0.08206

b. 62.36

c. 1.00

d. 22.4

121. How many moles of gas are there in a 45.0L container at 25°C and 500.0 mmHg?

a. 0.630

b. 6.11

c. 18.4

d. 1.21

122. How many moles of gas are there in a 50.0L container at 22.0°C and 825 torr?

a. 2.29 × 104

b. 1.70 × 103

c. 2.23

d. 0.603

123. Standard temperature and pressure (STP), in the context of gases, refers to

a. 298.15 K and 1 atm

b. 273.15 K and 1 atm

c. 298.15 K and 1 torr

d. 273.15 K and 1 pascal

124. Which one of the following gases would have the highest average molecular speed

at 25°C?

a. O2

b. N2

c. CO2

d. CH4

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125. A vessel contained N2, Ar, He, and Ne. The total pressure in the vessel was 987

torr. The partial pressures of nitrogen, argon, and helium were 44.0, 486, and 218 torr,

respectively. The partial pressure of neon in the vessel was ________ torr.

a. 42.4

b. 521

c. 19.4

d. 239

126. If the temperature of a fixed quantity of gas decreases and the pressure remains

unchanged,

a. its volume increases.

b. its volume is unchanged.

c. its volume decreases.

d. its density decreases.

127. If the pressure and temperature of a gas are held constant and some gas is added

to the container or some is allowed to escape, a change in which of the following can be

observed?

a. Kinetic energy

b. Volume

c. Rate of diffusion

d. Chemical properties

128. If the temperature of a container of gas remains constant, how could the

pressure of the gas increase?

a. The mass of the gas molecules increase.

b. The diffusion of the gas molecules increases.

c. The size of the container increases.

d. The number of gas molecules in the container increases.

129. If gas A has a molar mass greater than that of gas B and samples of each gas at

identical temperatures and pressures contain equal numbers of molecules, then

a. the volumes of gas A and gas B are equal.

b. the volume of gas A is greater than that of gas B.

c. the volume of gas B is greater than that of gas A.

d. their volumes are proportional to their molar masses.

130. The gas pressure inside a container decreases when

a. the number of gas molecules is increased.

b. the number of gas molecules is decreased.

c. the temperature in increased.

d. the number of molecules is increased and the temperature is increased.

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131. At a constant temperature, what would happen to the pressure of a gas if the

volume of the gas were doubled?

a. It would double.

b. It would decrease.

c. It would remain the same.

d. It would depend on the gas.

Thermochemistry

132. A solid forms when the average energy of a substance’s particles

a. increases.

b. decreases.

c. decreases then increases.

d. creates a random arrangement.

133. The energy of the particles in a solid is

a. higher than the energy of the particles in a gas.

b. high enough to allow the particles to interchange with other particles.

c. higher than the energy of the particles in a liquid.

d. lower than the energy of the particles in liquids and gases.

134. The compressibility of solids is generally

a. lower than the compressibility of liquids and gases.

b. higher than the compressibility of liquids only.

c. about equal to the compressibility of liquids and gases.

d. higher than the compressibility of gases only.

135. If the temperature and surface area of a liquid remain constant,

a. the liquid is not in equilibrium with its vapor.

b. no further evaporation occurs.

c. the rate of evaporation remains constant.

d. the rate of condensation is greater than the rate of evaporation.

136. If the temperature of a liquid-vapor system at equilibrium is reduced, the

a. concentration of the vapor will decrease.

b. rate of evaporation will increase.

c. equilibrium is unaffected.

d. percentage of liquid in the system will decrease.

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137. If the temperature of a liquid-vapor system at equilibrium increases, the new

equilibrium condition will

a. have a lower concentration of vapor.

b. have an increased vapor pressure.

c. not have equal rates of condensation and evaporation.

d. be larger in volume.

138. The energy transferred between samples of matter because of a difference in

their temperatures is called

a. heat.

b. thermochemistry.

c. chemical kinetics.

d. temperature.

139. Enthalpy change is the

a. pressure change of a system at constant temperature.

b. entropy change of a system at constant pressure.

c. temperature change of a system at constant pressure.

d. amount of energy absorbed or lost by a system as energy is the form of heat

during a process at constant pressure.

140. The quantity of energy released or absorbed as heat during a chemical reaction is

called the

a. temperature.

b. enthalpy of reaction.

c. entropy.

d. free energy.

141. In an endothermic reaction, the total energy at the beginning of the reaction is

a. greater than the total energy at the end of the reaction.

b. less than the total energy at the end of the reaction.

c. equal to the total energy at the end of the reaction.

d. None of the above.

142. For an exothermic reaction, the products

a. are at the same energy level as the reactants.

b. have no energy.

c. are at a lower energy level than the reactants.

d. are at a higher energy level than the reactants.

143. Ice melting is an example of a(n)

a. exothermic reaction.

b. negative entropic reaction.

c. endothermic reaction

d. catalyzed reaction.

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144. Entropy always increases as

a. the temperature of the system increases.

b. the temperature of the system decreases.

c. the enthalpy of the system decreases.

d. a chemical reaction proceeds.

145. Entropy in a system increases when

a. gases are diluted.

b. ions disperse in a solution.

c. the total moles of gaseous product exceed the total moles of gaseous reactant.

d. All of the above

146. As a general rule, the entropy of a solid is

a. less than that of a liquid.

b. more than that of a liquid.

c. more than that of a gas.

d. zero.

147. Which of the following systems has the lowest entropy?

a. Salt dissolved in a container of water

b. Sand mixed in a container of water

c. Oil floating in a container of water

d. A container of frozen water

148. Entropy decreases when

a. pressure decreases.

b. temperature decreases

c. the system is agitated.

d. temperature increases.

149. Which of the following has the highest entropy when produced in a reaction?

a. A solid

b. A liquid

c. A gas

d. An aqueous solution

150. Endothermic reactions feel

a. Cold

b. Hot

c. Lukewarm

d. Too hot to handle

151. What happens to the temperature during a phase change?

152. What happens to a reaction when the temperature is raised and why?

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Nuclear

Radiation comes from atoms with unstable nuclei- which usually means too many neutrons.

Fusion is nuclear change where two lighter elements “fuse” together to make a heavier element.

This is what happens in our Sun- 2 Hydrogen atoms fuse together to make a Helium atom.

Energy is always released.

Fission is nuclear change where a heavy element splits into lighter elements. This is used in

nuclear power and bombs. Energy is always released.

Types of radiation- alpha particles, beta particles, gamma rays, neutron emission

Half-life- is the time it takes for half of a sample to decay into another element

153. The half-life of an isotope is the time required for half the nuclei in a sample to

a. undergo radioactive decay.

b. undergo nuclear fission.

c. undergo nuclear fusion.

d. react chemically.

154. How many half-lives are required for three-fourths of the nuclei of one isotope in

a sample to decay?

a. 3/4

b. 3/2

c. 2

d. 3

155. Which statement is true about half-lives?

a. Different atoms of the same nuclide have different half-lives.

b. Each radioactive isotope has its own half-life.

c. All radioactive nuclides of an element have the same half-life.

d. All radioactive nuclides have the same half-life.

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Types of Reactions

Synthesis- a reaction in which two or more substances combine to form a new compound

Decomposition- a reaction in which a single compound breaks down to form two or more simpler

substances

Combustion- the oxidation reaction of an element or compound in which energy as heat is

released

Redox- any chemical change in which one species is oxidized (loses electrons) and another

species is reduced (gains electrons)

Neutralization- the reaction of the ions that characterize acids (hydronium ions) and the ions

that characterize bases (hydroxide ions) to form water molecules and a salt

156. In what kind of reaction do two or more substances combine to form a new

compound?

a. Decomposition reaction

b. Ionic reaction

c. Double-displacement reaction

d. Synthesis reaction

157. The equation AX A + X is the general equation for a

a. synthesis reaction.

b. decomposition reaction.

c. combustion reaction.

d. single-displacement reaction.

158. In what kind of reaction does one element replace a similar element in a

compound?

a. Displacement reaction

b. Combustion

c. Decomposition reaction

d. Ionic reaction

159. The equation AX + BY AY + BX is the general equation for a

a. synthesis reaction.

b. decomposition reaction.

c. single-displacement reaction.

d. double-displacement reaction.

160. In what kind of reaction does a single compound produce two or more simpler

substances?

a. Decomposition reaction

b. Synthesis reaction

c. Single-displacement reaction

d. Ionic reaction

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161. The equation A + X AX is the general equation for a(n)

a. combustion reaction.

b. ionic reaction.

c. synthesis reaction.

d. double-displacement reaction.

162. Write a balance chemical equation for the following reaction: iron plus copper (I)

nitrate yields iron (II) nitrate plus copper.

163. Write a balanced chemical equation for the synthesis of liquid phosphorus

trichloride, PCl3, from white phosphorus, P4, and chlorine gas.

164. Tell what type of chemical equation is represented by the following formula

equation. Then, balance the equation. C3H8 (g) + O2 (g) CO2 (g) + H2O (l)

165. Tell what type of chemical reaction is represented by the following formula

equation. Then, balance the equation. Ga (s) + S (s) Ga2S3 (s)

166. Tell what type of chemical reaction is represented by the following formula

equation. Then, balance the equation KBr (aq) + Mg(OH)2 (aq) KOH (aq) + MgBr2 (aq)

Stoichiometry

The Law of Conservation of Mass says that mass cannot be created or destroyed. In a

chemical reaction, this means that the number of each type of atom must be the same on each

side of the equation. It also means that you must have the same type of atoms on each side of

the reaction. The coefficients in the reaction tell you the constant ratio of the reactants and

products. Use the REP chart to help balance equations. For example, consider the reaction aX +

bY cZ. This tells us that a moles of X and b moles of Y will always react to produce c moles of

Z.

Molar mass is the mass in grams of 1 mole of a substance. Molar mass can be calculated by

adding the atomic masses of the atoms within a compound.

A mole ratio (a conversion factor) allows for the movement within a chemical reaction.

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167. Molar mass

a. is the mass in grams of one mole of a substance.

b. is numerically equal to the average atomic mass of the element.

c. Both (a) and (b).

d. Neither (a) nor (b).

168. What is the molar mass of magnesium chloride, MgCl2?

a. 46 amu

b. 59.76 amu

c. 95.21 amu

d. 106.35 amu

169. What is the molar mass of ethyl alcohol, C2H5OH?

a. 30.00 amu

b. 33.27 amu

c. 45.06 amu

d. 46.08 amu

170. What is the molar mass of (NH4)2SO4?

a. 114.09 amu

b. 118.34 amu

c. 128.06 amu

d. 132.16 amu

171. The word equation “solid carbon + oxygen gas carbon dioxide + energy”,

represents a chemical reaction because

a. the reaction releases energy.

b. CO2 has the chemical properties that differ from those of C and O.

c. the reaction absorbs energy.

d. CO2 is a gas and carbon is a crystal.

172. The complete balanced equation for the reaction between zinc hydroxide and

acetic acid is

a. ZnOH + CH3COOH ZnCH3COO + H2O

b. Zn(OH)2 + CH3COOH Zn + 2CO2 +3H2O

c. Zn(OH)2 + 2CH3COOH Zn(CH3COO)2 + 2H2O

d. Zn(OH)2 + 2CH3COOH Zn(CH3COO)2 + H2 + O2

173. What is the balanced equation for the combustion of sulfur?

a. S(s) + O2(g) SO(g)

b. S(s) + O2(g) SO2(g)

c. 2S(s) + 3O2(g) SO3(s)

d. S(s) + 2O2(g) SO42–(aq)

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174. Which equation is not balanced?

a. 2H2 + O2 2H2O

b. 4H2 + 2O2 4H2O

c. H2 + H2 + O2 H2O + H2O

d. 2H2 + O2 H2O

175. The Haber process for producing ammonia commercially is represented by the

equation N2(g) + 3H2(g) 2NH3(g). To completely convert 9.0 mol hydrogen gas to

ammonia gas, how many moles of nitrogen gas are required?

a. 1.0 mol

b. 2.0 mol

c. 3.0 mol

d. 6.0 mol

176. In the equation 2KClO3 2KCl + 3O2, how many moles of oxygen are produced

when 3.0 mol of KClO3 decompose completely?

a. 1.0 mol

b. 2.5 mol

c. 3.0 mol

d. 4.5 mol

177. For the reaction represented by the equation 2H2 + O2 2H2O, how many grams

of water are produced from 6.00 mol of hydrogen?

a. 2.00 g

b. 6.00 g

c. 54.0 g

d. 108 g

178. For the reaction represented by the equation 2Na + 2H2O 2NaOH + H2, how

many grams of sodium hydroxide are produced from 3.0 mol of sodium with an excess of

water?

a. 40.0 g

b. 80.0 g

c. 120 g

d. 240 g

179. How many moles are in 80 atoms of sodium?

a. 1.33 × 10-22 mol

b. 4.82 × 1025 mol

c. 2.37 × 10-37 mol

d. 3.95 × 10-20 mol

Equilibrium

Equilibrium- the state in which a chemical reaction and the reverse chemical reaction occur at

the same rate such that the concentrations of reactants and products do not change.

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180. At equilibrium,

a. all reactions have ceased.

b. only the forward reaction continues.

c. only the reverse reaction continues.

d. both the forward and reverse reactions continue.

181. If the system 2CO (g) + O2 (g) 2CO2 (g) has come to equilibrium and then the

volume is decreased,

a. [CO2] increases and [O2] decreases.

b. [CO2] increases and [O2] increases.

c. [CO2] decreases and [O2] decreases.

d. both [CO2] and [O2] remain the same.

182. If the pressure on the equilibrium system 2CO (g) + O2 (g) 2CO2 (g) is increased

a. the quantity of CO (g) increases.

b. the quantity of CO2 (g) decreases.

c. the quantity of CO2 (g) increases.

d. the quantities in the system do not change.

183. If the pressure on the equilibrium system N2 (g) + O2 (g) 2NO (g) decreases,

a. the quantity of N2 (g) decreases.

b. the quantity of NO (g) increases.

c. the quantity of NO (g) decreases.

d. the quantities in the system do not change.

184. If the temperature of the equilibrium system CH3OH (g) + 101 kJ CO (g) + 2H2

(g) increases,

a. [CH3OH] increases and [CO] decreases.

b. [CH3OH] decreases and [CO] increases.

c. [CH3OH] increases and [CO] increases.

d. the concentrations in the system do not change.

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