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Seniors Chem Review

Total questions: 128

Worksheet time: 32hrs 0mins

Name
Class
Date
1.

standard temperature and pressure

a)

the volume occupied by a mole of any gas at STP (22.4 L)

b)

the mass in amus of an atom of an element

c)

the number of representative particles in a mole of a substance

d)

the percent by mass of each element in a compound

e)

0°C and 101.3 kPa

2.

molar volume

a)

the volume occupied by a mole of any gas at STP (22.4 L)

b)

the mass in amus of an atom of an element

c)

the number of representative particles in a mole of a substance

d)

the percent by mass of each element in a compound

e)

0°C and 101.3 kPa

3.

atomic mass

a)

the volume occupied by a mole of any gas at STP (22.4 L)

b)

the mass in amus of an atom of an element

c)

the number of representative particles in a mole of a substance

d)

the percent by mass of each element in a compound

e)

0°C and 101.3 kPa

4.

Avogadro’s number

a)

the volume occupied by a mole of any gas at STP (22.4 L)

b)

the mass in amus of an atom of an element

c)

the number of representative particles in a mole of a substance

d)

the percent by mass of each element in a compound

e)

0°C and 101.3 kPa

5.

Avogadro’s number

a)

the volume occupied by a mole of any gas at STP (22.4 L)

b)

the mass in amus of an atom of an element

c)

the number of representative particles in a mole of a substance

d)

the percent by mass of each element in a compound

e)

0°C and 101.3 kPa

6.

percent composition

a)

the volume occupied by a mole of any gas at STP (22.4 L)

b)

the mass in amus of an atom of an element

c)

the number of representative particles in a mole of a substance

d)

the percent by mass of each element in a compound

e)

0°C and 101.3 kPa

7.

molar mass

a)

the mass of a mole of a substance

b)

the formula that gives the lowest whole-number ratio of the elements in a compound

c)

an atom, a formula unit, or a molecule

d)

equal volumes of gases at the same temperature and presssure contain equal numbers of particles

e)

6.02 x 1023 representative particles of a substance

8.

empirical formula

a)

the mass of a mole of a substance

b)

the formula that gives the lowest whole-number ratio of the elements in a compound

c)

an atom, a formula unit, or a molecule

d)

equal volumes of gases at the same temperature and presssure contain equal numbers of particles

e)

6.02 x 1023 representative particles of a substance

9.

representative particle

a)

the mass of a mole of a substance

b)

the formula that gives the lowest whole-number ratio of the elements in a compound

c)

an atom, a formula unit, or a molecule

d)

equal volumes of gases at the same temperature and presssure contain equal numbers of particles

e)

6.02 x 1023 representative particles of a substance

10.

Avogadro’s hypothesis

a)

the mass of a mole of a substance

b)

the formula that gives the lowest whole-number ratio of the elements in a compound

c)

an atom, a formula unit, or a molecule

d)

equal volumes of gases at the same temperature and presssure contain equal numbers of particles

e)

6.02 x 1023 representative particles of a substance

11.

mole (mol)

a)

the mass of a mole of a substance

b)

the formula that gives the lowest whole-number ratio of the elements in a compound

c)

an atom, a formula unit, or a molecule

d)

equal volumes of gases at the same temperature and presssure contain equal numbers of particles

e)

6.02 x1023 representative particles of a substance

12.

How many hydrogen atoms are in 4 molecules of isopropyl alcohol, C3H8O?

a)

4 x (6.02 x 1023)

b)

32

c)

4

d)

32 x (6.02 x 1023)

13.

The mass of a mole of NaCl is its:

a)

molar mass

b)

gram atomic mass.

c)

empirical formula.

d)

atomic mass.

14.

The chemical formula of aspirin is C9H8O4. What is the mass of 0.200 mol of aspirin?

a)

22.5 g

b)

5.4 g

c)

80 g

d)

36.0 g

15.

How many moles of SO3 are in 2.4 x 1024 molecules of SO3?

a)

0.25

b)

3.4 x 1022

c)

4.0

d)

2.9 x 10–23

16.

What is the volume (in liters at STP) of 2.50 mol of carbon monoxide?

a)

0.112 L

b)

3.10 L

c)

56.0 L

d)

8.96 L

17.

The molar mass of molecular oxygen (O2) is:

a)

equal to the mass of one mole of oxygen atoms.

b)

16.0 g.

c)

32.0 g.

d)

none of the above

18.

At STP, 1 mol each of hydrogen, oxygen, nitrogen, and fluorine:

a)

have different densities.

b)

occupy different volumes.

c)

are monatomic elements.

d)

contain twice Avogadro’s number of representative particles.

19.

a whole number that appears before a formula in an equation

a)

coefficient

b)

spectator ion

c)

combustion reaction

d)

reactant

e)

skeleton equation

20.

a particle not directly involved in a chemical reaction

a)

coefficient

b)

spectator ion

c)

combustion reaction

d)

reactant

e)

skeleton equation

21.

a reaction in which oxygen reacts with another substance, often producing heat or light

a)

coefficient

b)

spectator ion

c)

combustion reaction

d)

reactant

e)

skeleton equation

22.

a starting substance in a chemical reaction

a)

coefficient

b)

spectator ion

c)

combustion reaction

d)

reactant

e)

skeleton equation

23.

a chemical equation that does not indicate relative amounts of reactants and products

a)

coefficient

b)

spectator ion

c)

combustion reaction

d)

reactant

e)

skeleton equation

24.

an equation in which each side has the same number of atoms of each element

a)

balanced equation

b)

product

c)

decomposition reaction

d)

single-replacement reaction

25.

a new substance formed in a chemical reaction

a)

balanced equation

b)

product

c)

decomposition reaction

d)

single-replacement reaction

26.

a reaction in which a single compound is broken down into simpler substances

a)

balanced equation

b)

product

c)

decomposition reaction

d)

single-replacement reaction

27.

a reaction in which the atoms of one element replace the atoms of a second element

a)

balanced equation

b)

product

c)

decomposition reaction

d)

single-replacement reaction

28.

A catalyst is:

a)

a solid product of a reaction.

b)

one of the reactants in a single-replacement reaction.

c)

not used up in a reaction.

d)

the product of a combustion reaction.

29.

When the equation Fe + O2 ---> Fe2O3 is balanced, the coefficient for O2 is:

a)

4

b)

3

c)

2

d)

1

30.

The reaction Fe + O2 --> Fe2O3 is an example of a(n):

a)

aqueous reaction.

b)

combination reaction.

c)

single-replacement reaction.

d)

decomposition reaction.

31.

The equation H3PO4 + 3KOH ---> K3PO4 + 3H2O is an example of which type of reaction?

a)

double-replacement

b)

combination

c)

decomposition

d)

single-replacement

32.

In a double-replacement reaction:

a)

the reactants are usually a metal and a nonmetal.

b)

the reactants are generally two ionic compounds in aqueous solution.

c)

one of the reactants is often water.

d)

energy in the form of heat or light is always produced.

33.

In a double-replacement reaction:

a)

one of the products is always a gas.

b)

one of the products must be an element.

c)

positive ions are exchanged between two compounds.

d)

all of the above

34.

When the following equation is balanced, the coefficient in front of HCl is:

Ba(s) + HCl(aq) ---> BaCl2(aq) + H2(g)

a)

6

b)

3

c)

1

d)

2

35.

the amount of product formed when a reaction is carried out in the laboratory

a)

actual yield

b)

limiting reagent

c)

theoretical yield

d)

stoichiometry

36.

the calculated amount of product that might be formed during a reaction

a)

actual yield

b)

limiting reagent

c)

theoretical yield

d)

stoichiometry

37.

the reactant that determines the amount of product that can be formed in a reaction

a)

actual yield

b)

limiting reagent

c)

theoretical yield

d)

excess reagent

38.

the calculation of quantities in chemical equations

a)

actual yield

b)

limiting reagent

c)

theoretical yield

d)

stoichiometry

39.

the ratio of the actual yield to the theoretical yield, expressed as a percentage

a)

actual yield

b)

percent yield

c)

theoretical yield

d)

stoichiometry

40.

Which of these expressions is an incorrect interpretation of the balanced equation? 2S(s) + 3O2(g) → 2SO3(g)

a)

2 atoms S + 3 molecules O2 → 2 molecules SO3

b)

2 g S + 3 g O2 → 2 g SO3

c)

2 mol S + 3 mol O2 → 2 mol SO3

d)

none of the above

41.

In a chemical reaction, the mass of the products

a)

is less than the mass of the reactants.

b)

is greater than the mass of the reactants.

c)

is equal to the mass of the reactants.

d)

has no relationship to the mass of the reactants.

42.

How many liters of oxygen are required to react completely with 1.2 liters of hydrogen to form water?

2H2(g) + O2(g) → 2H2(g)

a)

1.2 L

b)

0.6 L

c)

2.4 L

d)

4.8 L

43.

How many liters of CO2(g) at STP are produced when 68.0 g of CaCO3(s) is heated according to the following equation?

CaCO3(s) → CaO(s) + CO2(g)

a)

0.679 L

b)

15.2 L

c)

68.0 L

d)

30.4 L

44.

A reaction that has been calculated to produce 60.0 g of CuCl2 actually produces 50.0 g of CuCl2. What is the percent yield?

a)

0.833%

b)

96.1%

c)

83.3%

d)

120%

45.

In any chemical reaction, the quantities that are conserved are

a)

the number of moles and the volumes.

b)

the number of molecules and the volumes.

c)

mass and number of atoms.

d)

mass and moles.

46.

3Cu(s) + 8HNO3(aq) → 3Cu(NO3)2(s) + 2NO(g) + 4H2O(l)


Calculate the number of moles of water produced when 3.3 mol of Cu(NO3)2 are formed in the reaction.

a)

4.4 mol

b)

6.6 mol

c)

4.9 mol

d)

8.8 mol

47.

3Cu(s) + 8HNO3(aq) → 3Cu(NO3)2(s) + 2NO(g) + 4H2O(l)


How many grams of Cu would be needed to react with 2.0 mol HNO3?

a)

95.3 g

b)

63.5 g

c)

47.6 g

d)

1.50 g

48.

describes a solid that lacks an ordered internal structure

a)

amorphous

b)

unit cell

c)

crystal

d)

normal boiling point

e)

atmospheric pressure

49.

the smallest group of particles within a crystal that retains the shape of the crystal

a)

amorphous

b)

unit cell

c)

crystal

d)

normal boiling point

e)

atmospheric pressure

50.

a solid in which the particles are arranged in an orderly, repeating, three-dimensional pattern

a)

amorphous

b)

unit cell

c)

crystal

d)

normal boiling point

e)

atmospheric pressure

51.

the temperature at which a liquid boils at a pressure of 101.3 kPa

a)

amorphous

b)

unit cell

c)

crystal

d)

normal boiling point

e)

atmospheric pressure

52.

the pressure resulting from the collision of particles in air with objects

a)

amorphous

b)

unit cell

c)

crystal

d)

normal boiling point

e)

atmospheric pressure

53.

the temperature at which a solid changes into a liquid

a)

melting point

b)

sublimation

c)

boiling point

d)

kinetic theory

54.

the change of a solid to a vapor without passing through the liquid state

a)

melting point

b)

sublimation

c)

boiling point

d)

kinetic theory

55.

the temperature at which the vapor pressure of a liquid is equal to the external pressure

a)

melting point

b)

sublimation

c)

boiling point

d)

kinetic theory

56.

states that the tiny particles in all forms of matter are in constant motion

a)

melting point

b)

sublimation

c)

boiling point

d)

kinetic theory

57.

one of two or more different molecular forms of the same element in the same physical state

a)

allotrope

b)

vapor pressure

c)

evaporation

d)

vacuum

58.

a measure of the force exerted by a gas above a liquid

a)

allotrope

b)

vapor pressure

c)

evaporation

d)

vacuum

59.

the conversion of a liquid to a gas or vapor at a temperature below the boiling point

a)

allotrope

b)

vapor pressure

c)

evaporation

d)

vacuum

60.

an empty space with no particles of matter

a)

allotrope

b)

vapor pressure

c)

evaporation

d)

vacuum

61.

The average kinetic energy of water molecules is greatest in

a)

steam at 200°C.

b)

liquid water at 90°C.

c)

liquid water at 373 K.

d)

ice at 0°C.

62.

According to the kinetic theory of gases,

a)

the particles in a gas move rapidly.

b)

the particles in a gas are relatively far apart.

c)

the particles in a gas move independently of each other.

d)

all are true.

63.

The temperature at which the motion of particles theoretically ceases is

a)

0°C.

b)

273°C.

c)

–273 K.

d)

0 K.

64.

The average kinetic energy of particles of a substance

a)

is not affected by the temperature of the substance.

b)

increases as the temperature of the substance decreases.

c)

is directly proportional to the temperature of a substance.

d)

is equal to 0.

65.

Which of these statements is not true, according to kinetic theory?

a)

There is no attraction between particles of a gas.

b)

Only particles of matter in the gaseous state are in constant motion.

c)

The particles of a gas collide with each other and with other objects.

d)

All collisions between particles of gas are perfectly elastic.

66.

Standard conditions when working with gases are defined as

a)

0 K and 101.3 kilopascals.

b)

0 K and 1 mm Hg.

c)

0°C and 101.3 kilopascals.

d)

0°C and 1 mm Hg.

67.

A phase diagram gives information on

a)

volumes of gases.

b)

conditions at which a substance exists as a solid, liquid, and gas.

c)

volumes of liquids and solids.

d)

changes in mass of solids, liquids, and gases.

68.

An increase in the temperature of a contained liquid

a)

causes the vapor pressure above the liquid to increase.

b)

decreases the vapor pressure above a liquid.

c)

causes fewer particles to escape the surface of the liquid.

d)

has no effect on the kinetic energy of the liquid.

69.

Water could be made to boil at 105°C by

a)

applying a great deal of energy.

b)

increasing the air pressure above the water.

c)

heating the water more gradually.

d)

decreasing the air pressure above the water.

70.

The escape of molecules from the surface of an uncontained liquid is

a)

boiling.

b)

sublimation.

c)

evaporation.

d)

condensation.

71.

8.31 (L x kPa) / (K x mol)

a)

ideal gas constant (R)

b)

Boyle’s law

c)

Dalton’s law of partial pressures

d)

ideal gas law

72.

For a given mass of gas at constant temperature, the volume of gas varies inversely with the pressure.

a)

ideal gas constant (R)

b)

Boyle’s law

c)

Dalton’s law of partial pressures

d)

ideal gas law

73.

At constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases.

a)

ideal gas constant (R)

b)

Boyle’s law

c)

Dalton’s law of partial pressures

d)

ideal gas law

74.

P × V = n × R × T

a)

ideal gas constant (R)

b)

Boyle’s law

c)

Dalton’s law of partial pressures

d)

ideal gas law

75.

P1 x V1 / T1 = P2 x V2 / T2

a)

combined gas law

b)

Charles’s law

c)

diffusion

d)

partial pressure

76.

The volume of a fixed mass of gas is directly proportional to the Kelvin temperature if the volume is kept constant.

a)

combined gas law

b)

Charles’s law

c)

diffusion

d)

partial pressure

77.

the contribution each gas in a mixture makes to the total pressure

a)

combined gas law

b)

Charles’s law

c)

diffusion

d)

partial pressure

78.

A gas tends to move to an area of lower concentration until the concentration is uniform throughout.

a)

combined gas law

b)

Charles’s law

c)

diffusion

d)

partial pressure

79.

As the temperature of a fixed volume of gas increases, the pressure will

a)

vary inversely.

b)

be unchanged.

c)

decrease.

d)

increase.

80.

A breathing mixture used by deep-sea divers contains helium, oxygen, and carbon dioxide. What is the partial pressure of oxygen at 101.3 kPa total pressure if PHe = 84.0 kPa and Pco2 = 0.10 kPa?

a)

10.3 kPa

b)

34.4 kPa

c)

17.2 kPa

d)

185.4 kPa

81.

Increasing the volume of a given amount of gas at constant temperature causes the pressure to decrease because

a)

the molecules are striking a larger area with the same force.

b)

there are fewer molecules.

c)

the molecules are moving more slowly.

d)

there are more molecules.

82.

When a container is filled with 3.00 mol of H2, 2.00 mol of O2, and 1.00 mol of N2, the pressure in the container is 465 kPa. The partial pressure of O2 is

a)

78 kPa.

b)

116 kPa.

c)

155 kPa.

d)

212 kPa.

83.

The volume of a gas is doubled while the temperature is held constant. The pressure of the gas

a)

remains unchanged.

b)

is reduced by one half.

c)

is doubled.

d)

depends on the kind of gas.

84.

The high surface tension of water is due to the:

a)

small size of water molecules.

b)

low mass of water molecules.

c)

hydrogen bonding between water molecules.

d)

covalent bonds in water molecules.

85.

Salts and other compounds that remove moisture from air are said to be:

a)

efflorescent.

b)

surfactant.

c)

colloidal.

d)

hygroscopic.

86.

The density of ice is less than the density of water because:

a)

ice has a lower molecular mass than water.

b)

the same mass occupies a smaller volume.

c)

the molecules are more closely packed.

d)

hydrogen bonding in ice produces an open framework.

87.

A solution is a mixture:

a)

from which the solute cannot be filtered.

b)

that is colloidal.

c)

that is heterogeneous.

d)

in which a solid solute is always dissolved in a liquid solvent.

88.

Which of the following is not an electrolyte?

a)

cane sugar(aq)

b)

HCl(aq)

c)

KCl(aq)

d)

(NH4)2SO4(aq)

89.

An electric current is conducted by:

a)

a solution of NaCl.

b)

a sugar solution.

c)

solid NaCl.

d)

solid sugar.

90.

Which of these would you expect to be soluble in the nonpolar solvent carbon disulfide, CS2?

a)

MgCl2

b)

CBr4

c)

CaCO3

d)

H2O

91.

Gelatin would best be classed as:

a)

a colloidal dispersion.

b)

a suspension.

c)

a heterogeneous mixture.

d)

an aqueous solution.

92.

A typical kind of emulsion is:

a)

muddy water.

b)

mayonnaise.

c)

sea water.

d)

smoke.

93.

When sodium chloride is mixed with water, it forms:

a)

a dispersion.

b)

an emulsion.

c)

a solution.

d)

a suspension.

94.

a solution containing the maximum amount of solute that can be dissolved at a given temperature

a)

saturated solution

b)

colligative properties

c)

miscible

d)

molarity

e)

unsaturated solution

95.

depend upon the number of particles of solute in solution

a)

saturated solution

b)

colligative properties

c)

miscible

d)

molarity

96.

describes liquids that dissolve in each other

a)

saturated solution

b)

colligative properties

c)

miscible

d)

molarity

e)

unsaturated solution

97.

the number of moles of solute dissolved in 1 L of solution

a)

saturated solution

b)

colligative properties

c)

miscible

d)

molarity

e)

unsaturated solution

98.

contains less solute than can theoretically be dissolved

a)

saturated solution

b)

colligative properties

c)

miscible

d)

molarity

e)

unsaturated solution

99.

describes liquids that are insoluble in each other

a)

immiscible

b)

concentrated solution

c)

supersaturated solution

d)

dilute solution

100.

a solution with a large amount of solute compared to solvent

a)

immiscible

b)

concentrated solution

c)

supersaturated solution

d)

dilute solution

101.

contains more solute than can theoretically be held at a given temperature

a)

immiscible

b)

concentrated solution

c)

supersaturated solution

d)

dilute solution

102.

contains only a small amount of the maximum amount of solute that can be dissolved at a given temperature

a)

immiscible

b)

concentrated solution

c)

supersaturated solution

d)

dilute solution

103.

Increasing the temperature of a liquid–solid solution will:

a)

a. always increase the rate at which a crystalline solute dissolves.

b)

b. often increase the amount of crystalline solute that dissolves.

c)

c. both a and b

d)

d. neither a nor b

104.

Which of the following operations usually makes a substance dissolve faster in a solvent?

a)

agitation

b)

crushing the substance to a powder

c)

raising the temperature

d)

all of the above

105.

To increase the solubility of a gas at constant temperature and 202 kPa pressure from 0.85 g/L to 5.1 g/L, the pressure would have to be increased to:

a)

1212 kPa.

b)

606 kPa.

c)

505 kPa.

d)

17.2 kPa.

106.

If the pressure of a gas above a liquid is decreased (at constant temperature), the solubility of the gas in the liquid:

a)

remains unchanged.

b)

increases.

c)

decreases.

d)

would change but in an unpredictable direction.

107.

How many mL of alcohol are in 240 mL of 95.0% (v/v) alcohol solution?

a)

12 mL

b)

145 mL

c)

228 mL

d)

142 mL

108.

If more solvent is added to a solution:

a)

the molarity decreases.

b)

the solution becomes less dilute.

c)

the percent (v/v) increases.

d)

all of the above

109.

What is the molarity of a 200-mL solution in which 0.2 mole of sodium bromide is dissolved?

a)

0.20M

b)

0.40M

c)

1.0M

d)

4.0M

110.

Which of the following is not a colligative property of a solution?

a)

boiling-point elevation

b)

vapor-pressure lowering

c)

solubility

d)

freezing-point depression

111.

the amount of heat required to change the temperature of an object by exactly 1°C

a)

heat capacity

b)

law of conservation of energy

c)

heat of reaction

d)

energy

112.

In any chemical or physical process, energy is neither created nor destroyed.

a)

heat capacity

b)

law of conservation of energy

c)

heat of reaction

d)

energy

113.

The heat absorbed or released by a chemical reaction

a)

heat capacity

b)

law of conservation of energy

c)

heat of reaction

d)

energy

114.

the capacity to do work or to supply heat

a)

heat capacity

b)

law of conservation of energy

c)

heat of reaction

d)

energy

115.

energy that always flows from a warmer object to a cooler object

a)

heat

b)

standard heat of formation

c)

calorie

d)

exothermic process

e)

specific heat

116.

the quantity of heat that raises the temperature of 1 g of pure water 1°C

a)

heat

b)

standard heat of formation

c)

calorie

d)

exothermic process

e)

specific heat

117.

the change in enthalpy that accompanies the formation of 1 mole of a compound from its elements, with all substances in their standard states at 25°C

a)

heat

b)

standard heat of formation

c)

calorie

d)

exothermic process

e)

specific heat

118.

a process that loses heat to the surroundings

a)

heat

b)

standard heat of formation

c)

calorie

d)

exothermic process

e)

specific heat

119.

the amount of heat required to raise the temperature of 1 gram of a substance 1°C

a)

heat

b)

standard heat of formation

c)

calorie

d)

exothermic process

e)

specific heat

120.

The SI unit of energy is

a)

heat capacity.

b)

joule.

c)

calorie.

d)

enthalpy.

121.

How many calories are required to raise the temperature of 75.0 g of water from 20°C to 50°C?

a)

1.50 × 103 cal

b)

3750 cal

c)

2250 cal

d)

75.0 cal

122.

The temperature of a 6.0-g sample of glass changed from 20°C to 45°C when it absorbed 550 J of heat. What is the specific heat of this glass sample?

a)

3.7 J/g•°C

b)

2300 J/g•°C

c)

0.27 J/g•°C

d)

130 J/g•°C

123.

The enthalpy of a system is the same as its:

a)

specific heat.

b)

heat of combustion.

c)

heat of reaction.

d)

heat content.

124.

When your body is warmed by an electric blanket during the winter, this process is said to be

a)

endothermic.

b)

isothermic.

c)

exothermic.

d)

none of the above

125.

A student mixes two water solutions with an initial temperature of 25.0°C to form a final solution with a mass of 65.0 g at 30.0°C. What is the heat change, in kJ, for this reaction?

a)

325 kJ

b)

1.36 kJ

c)

272 kJ

d)

1.95 kJ

126.

Given the equation 2Mg(s) + O2(g) ---> 2MgO(s) + 72.3 kJ, which of the following is true?

a)

The reaction is endothermic.

b)

delta H = +72.3 kJ

c)

delta H = –72.3 kJ

d)

The reaction absorbs heat.

127.

Given the equation Si(s) + 2Cl2(g) ®SiCl2(g) ---> SiCl2(l) + 687 kJ, how much heat is produced when 106 g of Cl2 react?

a)

7.28 × 104 kJ

b)

360 kJ

c)

513 kJ

d)

180 kJ

128.

How much heat, in kJ, is required to melt 54.0 g of ice at 0°C into water at 0°C if deltaHfus for water = 6.01 kJ/mol?

a)

0.111 kJ

b)

18.0 kJ

c)

325 kJ

d)

8.99 kJ

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