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Final Exam Review

Total questions: 177

Worksheet time: 2hrs 4mins

Name
Class
Date
1.

What is a reactant in a chemical reaction?

a)

The new substance formed by a chemical reaction

b)

The starting substance of a chemical reaction

c)

The substance that speeds up a chemical reaction

d)

The substance that stops a chemical reaction

2.

What is a product in a chemical reaction?

a)

The starting substance of a chemical reaction

b)

The substance that speeds up a chemical reaction

c)

The new substance formed by a chemical reaction

d)

The substance that stops a chemical reaction

3.

What is a balanced chemical equation?

a)

An equation with more reactants than products

b)

An equation with more products than reactants

c)

An equation with the same number of each type of atom on both sides

d)

An equation with different numbers of atoms on each side

4.

What does the symbol (s) represent in a chemical equation?

a)

Solid

b)

Liquid

c)

Gas

d)

Aqueous solution

5.

What does the symbol (aq) represent in a chemical equation?

a)

Solid

b)

Liquid

c)

Gas

d)

Dissolved in water

6.

What is one of the signals that a chemical reaction has occurred?

a)

No change in color

b)

Solid formation

c)

No change in temperature

d)

No change in state

7.

What must be checked to ensure a chemical equation is balanced?

a)

The color of the reactants

b)

The state of matter

c)

The number of atoms on both sides

d)

The temperature of the reaction

8.

What does a balanced chemical equation ensure?

a)

Energy will transfer from reactants to products

b)

Equal volume of reactants and products

c)

Equal number of molecules

d)

Equal number of atoms of each type

9.

What is not a part of the systematic approach to balancing equations?

a)

Writing unbalanced equations

b)

Changing subscripts

c)

Balancing by trial and error

d)

Checking atom balance

10.

Why is it important to balance chemical equations?

a)

To satisfy the law of conservation of mass

b)

To make the equation look neat

c)

To change the properties of reactants

d)

To increase the reaction speed

11.

What does the coefficient in a chemical equation represent?

a)

The speed of the reaction

b)

The temperature at which the reaction occurs

c)

The relative number of molecules

d)

The type of chemical bond

12.

Which of the following is a key idea in balancing chemical equations?

a)

Use only even numbers for coefficients

b)

Balance by changing the chemical formulas

c)

Balance by trial and error

d)

Use only odd numbers for coefficients

13.

What is the role of coefficients in a chemical equation?

a)

To indicate the phase of the substance

b)

To show the relative amounts of reactants and products

c)

To change the chemical identity

d)

To denote the temperature of the reaction

14.

What is the purpose of writing the formulas of reactants and products in a chemical equation?

a)

To change the chemical properties

b)

To create an unbalanced equation

c)

To start the balancing process

d)

To remove unnecessary elements

15.
Bubbles form when two or more substances are combined ...
a)
Chemical Change
b)
Physical Change
16.

​ (a)   Fe(s)Fe_{\left(s\right)} + ​ ​ (b)   Cl2(g)Cl_{2\left(g\right)} →\rightarrow ​ (c)   FeCl3(s)FeCl_{3\left(s\right)}

Balance the equation by adding coefficients.

Choose from the below words
2
3
1
17.

Balance this equation:

​ (a)   Fe + ​ (b)   O2 --> ​ (c)   Fe2O3

Choose from the below words
4
3
2
1
0
18.

In 3P2O5, there are ​ (a)   phosphorus (P) atoms and ​ (b)   oxygen (O) atoms.

Choose from the below words
6
15
8
3
12
21
19.

Number of C: ​ (a)  

Number of H: ​ (b)  

Number of O: ​ (c)  

Choose from the below words
4
6
8
2
3
1
5
7
9
10
20.

4CF2Cl3

a)

carbon

1.

4

b)

fluorine

2.

8

c)

chlorine

3.

12

21.

Balance the following reaction:

​ (a)   C + ​ (b)   H2 → ​ (c)   C3H8

Choose from the below words
3
4
1
2
5
22.

Say how many atoms of oxygen are in each of the following.

a)

2

1.

NaNO2NaNO_2

b)

4

2.

2NaNO22NaNO_2

c)

8

3.

2Mg(NO2)22Mg\left(NO_2\right)_2

23.

Two clear chemicals combine. The solution turns yellow and gets cold. What evidence suggests that a chemical change took place?

a)

There is no evidence.

b)

The solution turned yellow.

c)

The solution got cold.

d)

The solution both turned yellow and got cold.

24.

Which of the following is NOT an indicator (evidence) of a chemical reaction?

a)

Precipitate formation

b)

Light being produced

c)

Change in state of matter

d)

Change in temperature

25.

How are fireworks an example of chemical reactions?

a)

The chemicals do not change as they explode.

b)

They are made of chemicals.

c)

Heat, light, and sound are given off.

d)

Fireworks are not examples of chemical reactions.

26.

What is the force that holds two or more atoms together and makes them function as a unit?

a)

Bond energy

b)

Ionic bonding

c)

Bond

d)

Covalent bonding

27.

What are some of the general properties of gases that distinguish them from liquids and solids?

a)

Gases have a definite shape and volume.

b)

Gases are incompressible.

c)

Gases have high density.

d)

Gases have no definite shape.

e)

Gases are compressible.

28.

What is the SI unit of pressure?

a)

Pascal

b)

Bar

c)

PSI

d)

Torr

29.

Convert 1.20 atm to units of mm Hg, torr, and pascals.

a)

1.20 atm is equivalent to 912 mm Hg, 912 torr, and 121,000 pascals.

b)

1.20 atm is equivalent to 900 mm Hg, 900 torr, and 120,000 pascals.

c)

1.20 atm is equivalent to 920 mm Hg, 920 torr, and 122,000 pascals.

d)

1.20 atm is equivalent to 930 mm Hg, 930 torr, and 123,000 pascals.

30.

Boyle’s law for gases tells us that:

a)

The pressure of a gas is inversely proportional to its volume at constant temperature.

b)

The pressure of a gas is directly proportional to its volume at constant temperature.

c)

The pressure of a gas is inversely proportional to its temperature at constant volume.

d)

The pressure of a gas is directly proportional to its temperature at constant volume.

31.

What does Charles’s law tell us about how the volume of a gas sample varies as the temperature of the sample is changed?

a)

The volume of a gas is directly proportional to its temperature.

b)

The volume of a gas is inversely proportional to its temperature.

c)

The volume of a gas remains constant as temperature changes.

d)

The volume of a gas is independent of its temperature.

32.

What temperature scale is defined with its lowest point as the absolute zero of temperature?

a)

Kelvin scale

b)

Celsius scale

c)

Fahrenheit scale

d)

Rankine scale

33.

What is absolute zero in Celsius degrees?

a)

-273.15°C

b)

0°C

c)

-459.67°C

d)

273°C

34.

What does Avogadro’s law tell us about the relationship between the volume of a sample of gas and the number of molecules the gas contains?

a)

Avogadro's law states that the volume of a gas is directly proportional to the number of molecules, at constant temperature and pressure.

b)

Avogadro's law states that the volume of a gas is inversely proportional to the number of molecules, at constant temperature and pressure.

c)

Avogadro's law states that the volume of a gas is independent of the number of molecules, at constant temperature and pressure.

d)

Avogadro's law states that the volume of a gas is directly proportional to the temperature, at constant pressure.

35.

What equation would you use to solve this?

A sample of gas in a 10.0-L container exerts a pressure of 565 mm Hg. Calculate the pressure exerted by the gas if the volume is changed to 15.0 L at constant temperature.

a)

b)

c)

d)

36.

A sample of gas in a 10.0-L container exerts a pressure of 565 mm Hg. Calculate the pressure exerted by the gas if the volume is changed to 15.0 L at constant temperature.

(a)   mm Hg

37.

What equation would you use to solve this?

A sample of gas in a 5.00-L container at 35.0 °C is heated at constant pressure to a temperature of 70.0 °C at constant pressure. Determine the volume of the heated gas.

a)

b)

c)

d)

38.

A sample of gas in a 5.00-L container at 35.0 °C is heated at constant pressure to a temperature of 70.0 °C at constant pressure. Determine the volume of the heated gas.

_. (a)   L

39.

What equation would you use to solve this?

A sample of gas at 24 °C occupies a volume of 3.45 L and exerts a pressure of 2.10 atm. The gas is cooled to –12 °C, and the pressure is increased to 5.20 atm. Determine the new volume occupied by the gas.

a)

b)

c)

d)

40.

A sample of gas at 24 °C occupies a volume of 3.45 L and exerts a pressure of 2.10 atm. The gas is cooled to –12 °C, and the pressure is increased to 5.20 atm. Determine the new volume occupied by the gas.

_. (a)   L

41.

What equation would you use to solve this?

A 4.50-mol sample of a gas occupies a volume of 34.6 L at a particular temperature and pressure. What volume does 2.50 mol of the gas occupy at these same conditions of pressure and temperature?

a)

b)

c)

d)

42.

A 4.50-mol sample of a gas occupies a volume of 34.6 L at a particular temperature and pressure. What volume does 2.50 mol of the gas occupy at these same conditions of pressure and temperature?

(a)   ._ L

43.

What equation would you START WITH to solve this?

What mass of helium gas exerts a pressure of 1.20 atm in a volume of 5.40 L at a temperature of 27 °C?

a)

b)

c)

d)

44.

What mass of helium gas exerts a pressure of 1.20 atm in a volume of 5.40 L at a temperature of 27 °C?

_. (a)   g

45.

A 2.50-g sample of neon gas is added to a 5.00-g sample of argon gas in a 10.0-L container at 23 °C. Calculate the total pressure of the mixture.

_. (a)   atm

46.

A sample of oxygen gas is collected over water at 27 °C. The total pressure is 0.95 atm, and the water vapor pressure at 27 °C is 26.7 torr. Determine the partial pressure of the oxygen gas collected.

(a)   torr

47.

What does a barometer measure?

a)

Temperature

b)

Humidity

c)

Atmospheric pressure

d)

Volume

48.

What is another name for torr?

a)

Pascal

(Pa)

b)

Millimeters of mercury

(mm Hg)

c)

Atmosphere

(atm)

d)

Newton per square meter

(N/m2)

49.

How many mm Hg is equivalent to one standard atmosphere?

a)

500 mm Hg

b)

760 mm Hg

c)

990 mm Hg

d)

800 mm Hg

50.

What is the SI unit of pressure?

(equal to one newton per square meter)

a)

Atmosphere

b)

Torr

c)

Pascal

d)

Bar

51.

According to Boyle’s Law, what happens to the pressure of a gas when the volume increases at constant temperature?

a)

Pressure increases

b)

Pressure decreases

c)

Pressure remains the same

d)

Pressure doubles

52.

What temperature is defined as absolute zero?

a)

0 degrees Celsius

b)

–273 degrees Celsius

c)

–100 degrees Celsius

d)

100 degrees Celsius

53.

Which law states that the volume of a gas is directly related to its temperature at constant pressure?

a)

Boyle’s Law

b)

Charles’s Law

c)

Avogadro’s Law

d)

Dalton’s Law

54.

What does Avogadro’s Law state?

a)

Volume is inversely proportional to pressure

b)

Volume is directly proportional to temperature

c)

Equal volumes of gases contain the same number of particles

d)

Total pressure is the sum of partial pressures

55.

What is the universal gas constant (R) in L·atm/K·mol?

a)

8.314

b)

0.08206

c)

101.325

d)

760

56.

In the ideal gas law (PV = nRT), what does 'n' represent?

a)

Pressure

b)

Volume

c)

Moles of gas

d)

Temperature

57.

Ptotal = P1 + P2 + P3 + ...

a)

Dalton's Law of Partial Pressures

b)

Boyle's Law

c)

Combined Gas Law

d)

Avogadro's Law

58.

What is an ideal gas?

a)

A gas that does not exist

b)

A gas that obeys PV = nRT

c)

A gas at high pressure

d)

A gas at low temperature

59.

What is the combined gas law?

a)

PV = nRT

b)

P1V1/T1 = P2V2/T2

c)

P + V = T

d)

V = bT

60.

What is partial pressure?

a)

Total pressure in a container

b)

The pressure exerted by a single gas in a mixture

c)

The pressure when gas volumes are equal

d)

None of the above

61.

What does Dalton’s Law of Partial Pressures state?

a)

Total pressure is the sum of all gas volumes

b)

Total pressure is the sum of partial pressures of individual gases

c)

Total pressure equals pressure of one gas

d)

Gas volume is independent of temperature

62.

What is the molar volume of an ideal gas at standard temperature and pressure?

a)

1 liter

b)

22.42 liters

c)

10 liters

d)

8.24 liters

63.

What conditions define standard temperature and pressure (STP)?

a)

0 degrees Celsius and 1 atm

b)

0 Kelvin and 1 atm

c)

0 degrees Celsius and 760 mm Hg

d)

100 degrees Celsius and 1 atm

64.

What does the kinetic molecular theory assume about gases?

a)

Gases are composed of large particles

b)

Gases are always at high pressure

c)

Gases consist of tiny particles in constant motion

d)

Gases do not exert pressure

65.

Which of the following is true according to Boyle's law?

a)

Volume increases as pressure increases

b)

Pressure decreases as volume increases

c)

Temperature and volume are inversely related

d)

Pressure and temperature are independent

66.

What is the relationship described by Charles's law?

a)

Pressure is inversely proportional to volume

b)

Volume is directly proportional to temperature

c)

Pressure is directly proportional to volume

d)

Volume is independent of temperature

67.

Which law can be expressed as PV = k?

a)

Ideal Gas Law

b)

Boyle's Law

c)

Charles's Law

d)

Dalton’s Law

68.

How does the volume of a gas change when temperature increases at constant pressure?

a)

Volume decreases

b)

Volume stays the same

c)

Volume increases

d)

Volume doubles

69.

What happens to the behavior of real gases at high pressures and low temperatures?

a)

They behave ideally

b)

They do not behave ideally

c)

They become solids

d)

They expand indefinitely

70.

According to the combined gas law, if the temperature of a gas decreases while keeping the volume constant, what happens to the pressure?

a)

Pressure increases

b)

Pressure decreases

c)

Pressure remains the same

d)

Pressure cannot be determined

71.

In the ideal gas equation, which unit should T be in?

a)

Torr

b)

Pascal

c)

Celsius

d)

Kelvin

e)

g/mol

72.

Categorize the densities of solids liquids and gases

Categorize the following

Solid

Liquid

Gas

More Dense
Less Dense
73.

Categorize the compressibility of solids liquids and gases

Categorize the following

Solid

Liquid

Gas

Not Compressible
Compressible
74.

Categorize the characteristics of water

Categorize the following

colorless

odorless

freezes at 0°C

boils at 100°C

tasteless

bluish color

acidic taste

chlorine smell

boils at 100 K

freezes at 0 K

Water
Not Water
75.

Which is greater?

(if moles and pressure are constant)

a)

energy needed to boil water

b)

energy needed to melt ice

76.

Choose which would be 'changes in state'.

a)

liquid to gas

b)

gas to solid

c)

solid to liquid

d)

atom to molecule

e)

low pressure gas to

high pressure gas

77.

In order to melt or vaporize a substance, which forces might need to be overcome?

a)

London dispersion forces

b)

Dipole-Dipole forces

c)

Hydrogen bonding

d)

Covalent bonding

e)

Ionic Bonding

78.

The forces that must be overcome to vaporize a liquid are

a)

Intermolecular

b)

Intramolecular

79.

Define molar heat of fusion and molar heat of vaporization. Choose the correct option.

a)

Molar heat of fusion is the energy per mole required to change a substance from solid to liquid at constant temperature, and molar heat of vaporization is the energy per mole required to change it from liquid to gas at constant temperature.

b)

Molar heat of fusion is the energy per mole required to raise the temperature of a solid, and molar heat of vaporization is the energy per mole required to vaporize a gas.

c)

Molar heat of fusion and vaporization both refer to the energy changes during chemical reactions.

d)

Molar heat of fusion is the energy released during solidification, and molar heat of vaporization is the energy absorbed during condensation.

80.

The heat of fusion of aluminum is 3.95 kJ/g. Determine the molar heat of fusion of aluminum, given its molar mass is approximately 26.98 g/mol.

a)

107 kJ/mol

b)

96.5 kJ/mol

c)

30.93 kJ/mol

d)

6.83 kJ/mol

81.

Although London forces exist among all molecules, for what type of molecule are they the only major intermolecular force?

a)

Nonpolar molecules

b)

Polar molecules

c)

Covalent molecules

d)

Ionic molecules

e)

Metals

82.

Equilibrium vapor pressure of a liquid is defined as the pressure exerted by its vapor in equilibrium with the liquid. How is it related to the strength of intermolecular forces?

a)

It is higher for liquids with weaker intermolecular forces.

b)

It is higher for liquids with stronger intermolecular forces.

c)

It is independent of the intermolecular forces.

d)

It depends solely on temperature and not on intermolecular forces.

83.

What is the vapor pressure of water at 100.0 °C ?

a)

1 atm

b)

0.5 atm

c)

2 atm

d)

it depends on the elevation

84.

The term crystalline solid is defined as:

a)

A solid in which particles are arranged in a regular pattern.

b)

A solid with no fixed shape or volume.

c)

A solid where the particles are arranged randomly throughout.

d)

A solid that melts at a very high temperature.

85.

Which two are true of ALL alloys?

a)

metallic properties

b)

composed of two metals

c)

contain a mixture of elements

d)

corrosion resistant

86.

A sample of boiling water remains at the same temperature even though heat is continually added to it.

Where does the heat energy go?

a)

The water absorbs heat to increase its potential energy.

b)

The heat energy is used to change the water from liquid to gas.

c)

The heat energy decreases the vapor pressure inside the liquid.

d)

The water loses energy due to evaporation.

87.

boiling point ​ (a)   with an increase in altitude because atmospheric pressure ​ (b)  

Choose from the below words
decreases
increases
stays the same
88.

How do the strengths of dipole–dipole forces compare with the strengths of

typical covalent bonds?

a)

Dipole–dipole forces are significantly weaker than covalent bonds.

b)

It depends on the atoms or molecules involved

c)

Dipole–dipole forces are significantly stronger than covalent bonds.

d)

Dipole–dipole forces and covalent bonds have comparable strengths.

89.

Which condition must exist for hydrogen bonding to happen?

a)

hydrogen bonded to a very electronegative atom

b)

hydrogen is bonded to carbon

c)

any polar covalent bond is present

d)

two hydrogens must be present in each molecule

90.

The boiling point of a liquid is related to the atmospheric pressure because

a)

Increased atmospheric pressure requires higher energy to vaporize the liquid.

b)

Decreased atmospheric pressure requires higher energy to vaporize the liquid.

c)

Increased atmospheric pressure lowers the energy needed for vaporization.

d)

Atmospheric pressure does not affect the boiling point.

91.

Which of the following pair of crystalline solid and the interparticle forces holding its particles together is correctly matched?

a)

NaCl – Ionic solid, held together by electrostatic forces

b)

Diamond – Metallic solid, held together by the sea of electrons

c)

Copper – Covalent network solid, held together by directional covalent bonds

d)

Dry Ice (CO2) – Ionic solid, held together by ionic bonds

92.

Categorize the alloys

Categorize the following

brass

steel

Substitutional
Interstitial
93.

What instrument would be best to use in an experiment to demonstrate vapor pressure?

a)

manometer

b)

barometer

c)

pressure cooker

d)

anemometer

94.

Define a solution.

a)

A homogeneous mixture of two or more substances.

b)

A heterogeneous mixture of two or more substances.

c)

A pure substance.

d)

A compound.

95.

How does a molecular solid such as sugar dissolve in water?

a)

portions of sugar molecules are attracted to water molecules

b)

sugar molecules form covalent bonds with water molecules

c)

sugar molecules form ionic bonds with water molecules

d)

sugar molecules split into separate ions to dissolve in water

96.

How does an ionic solute such as NaCl dissolve in water?

a)

portions of NaCl crystals for hydrogen bonds with water molecules

b)

NaCl crystals form covalent bonds with water molecules

c)

NaCl crystals form ionic bonds with water molecules

d)

NaCl crystals split into separate ions to dissolve in water

97.

What are three ways to increase the rate of dissolution of a solute?

a)

heat the solution

b)

stir the solution

c)

increase the surface area of the solute

d)

increase the number of solute particles

98.

Match the following statements correctly with a saturated solution, an unsaturated solution, and a supersaturated solution?

a)

contains as much solute as can dissolve at a

particular temperature

1.

saturated solution

b)

contains less solute

than can dissolve at a particular temperature

2.

unsaturated solution

c)

more solute is dissolved than is normally possible at a particular

temperature

3.

supersaturated solution

99.

Match the terms to the definitions

a)

(mass of solute) × 100%

(mass of solution)

1.

Percent by mass

(or mass percent)

b)

the number of moles of solute per liter of solution

2.

Molarity

c)

the number of equivalents per liter of solution

3.

Normality

100.

One equivalent of an acid can furnish ​ (a)   of ​ (b)   ions.

One equivalent of a base can furnish ​ (c)   of​ (d)   ions.

Choose from the below words
one mole
H+
OH–
positive
negative
one kilogram
one gram
101.

The normality of a solution is defined as

a)

equivalents of solute

liter of solution

b)

equivalents of solvent

liter of solution

c)

moles of solute

liter of solution

d)

moles of solvent

liter of solution

102.

A 12.5-g sample of glucose (C₆H₁₂O₆) is dissolved in 225 g of water. Calculate the percent by mass of glucose in the solution.

a)

4.8%

b)

5.3%

c)

6.3%

d)

5.8%

103.

A chemist prepares some standard solutions for use in the lab using 500.0-mL volumetric flasks to contain the solutions. If the following masses of solutes are used, calculate the resulting molarity of a solution with 4.865 g NaCl. (The molar mass of NaCl is 58.44 g/mol)

a)

0.167 M

b)

0.132 M

c)

0.244 M

d)

0.929 M

104.

Suppose that each of the following solutions is diluted by adding the indicated amount of water. Calculate the new concentrations of

255 mL of 3.02 M HCl with 375 mL water added

a)

1.22 M HCl

b)

2.02 M HCl

c)

0.98 M HCl

d)

1.50 M HCl

105.

Calculate the volume (in milliliters) of

0.271 M HCl

that would be required to neutralize

36.2 mL of 0.259 M NaOH

a)

34.6 mL

b)

36.2 mL

c)

33.0 mL

d)

35.0 mL

106.

What volume of

0.242 M H2SO4

can furnish the same number of moles of H+ ions as

41.5 mL of 0.118 M HCl

a)

10.1 mL

b)

9.45 mL

c)

20.3 mL

d)

5.05 mL

107.

Calculate the normality of the following solution: 0.204 M HCl

a)

0.204 N

b)

0.102 N

c)

0.408 N

d)

0.306 N

108.

Calculate the normality of the following solution: 0.328 M H2SO4

a)

0.656 N

b)

0.328 N

c)

1.312 N

d)

0.164 N

109.

What volume of 0.10 M Ba(NO3)2 solution is required to react completely with 100.0 mL of 0.50 M NaCl solution to form BaCl2(s)?

a)

250.0 mL

b)

200.0 mL

c)

500.0 mL

d)

100.0 mL

110.

The term colligative property refers to a property of a solution that depends only on the

a)

number of solute particles

b)

chemical nature of the solute

c)

solvent characteristics

d)

the number and type of solute particles

111.

Define a solution.

a)

A homogeneous mixture of two or more substances.

b)

A heterogeneous mixture of two or more substances.

c)

A pure substance.

d)

A compound.

112.

How does a molecular solid such as sugar dissolve in water?

a)

portions of sugar molecules are attracted to water molecules

b)

sugar molecules form covalent bonds with water molecules

c)

sugar molecules form ionic bonds with water molecules

d)

sugar molecules split into separate ions to dissolve in water

113.

How does an ionic solute such as NaCl dissolve in water?

a)

portions of NaCl crystals for hydrogen bonds with water molecules

b)

NaCl crystals form covalent bonds with water molecules

c)

NaCl crystals form ionic bonds with water molecules

d)

NaCl crystals split into separate ions to dissolve in water

114.

What are three ways to increase the rate of dissolution of a solute?

a)

heat the solution

b)

stir the solution

c)

increase the surface area of the solute

d)

increase the number of solute particles

115.

Match the following statements correctly with a saturated solution, an unsaturated solution, and a supersaturated solution?

a)

contains as much solute as can dissolve at a

particular temperature

1.

saturated solution

b)

contains less solute

than can dissolve at a particular temperature

2.

unsaturated solution

c)

more solute is dissolved than is normally possible at a particular

temperature

3.

supersaturated solution

116.

Match the terms to the definitions

a)

(mass of solute) × 100%

(mass of solution)

1.

Percent by mass

(or mass percent)

b)

the number of moles of solute per liter of solution

2.

Molarity

c)

the number of equivalents per liter of solution

3.

Normality

117.

One equivalent of an acid can furnish ​ (a)   of ​ (b)   ions.

One equivalent of a base can furnish ​ (c)   of​ (d)   ions.

Choose from the below words
one mole
H+
OH–
positive
negative
one kilogram
one gram
118.

The normality of a solution is defined as

a)

equivalents of solute

liter of solution

b)

equivalents of solvent

liter of solution

c)

moles of solute

liter of solution

d)

moles of solvent

liter of solution

119.

A 12.5-g sample of glucose (C₆H₁₂O₆) is dissolved in 225 g of water. Calculate the percent by mass of glucose in the solution.

a)

4.8%

b)

5.3%

c)

6.3%

d)

5.8%

120.

A chemist prepares some standard solutions for use in the lab using 500.0-mL volumetric flasks to contain the solutions. If the following masses of solutes are used, calculate the resulting molarity of a solution with 4.865 g NaCl. (The molar mass of NaCl is 58.44 g/mol)

a)

0.167 M

b)

0.132 M

c)

0.244 M

d)

0.929 M

121.

Suppose that each of the following solutions is diluted by adding the indicated amount of water. Calculate the new concentrations of

255 mL of 3.02 M HCl with 375 mL water added

a)

1.22 M HCl

b)

2.02 M HCl

c)

0.98 M HCl

d)

1.50 M HCl

122.

Calculate the volume (in milliliters) of

0.271 M HCl

that would be required to neutralize

36.2 mL of 0.259 M NaOH

a)

34.6 mL

b)

36.2 mL

c)

33.0 mL

d)

35.0 mL

123.

What volume of

0.242 M H2SO4

can furnish the same number of moles of H+ ions as

41.5 mL of 0.118 M HCl

a)

10.1 mL

b)

9.45 mL

c)

20.3 mL

d)

5.05 mL

124.

Calculate the normality of the following solution: 0.204 M HCl

a)

0.204 N

b)

0.102 N

c)

0.408 N

d)

0.306 N

125.

Calculate the normality of the following solution: 0.328 M H2SO4

a)

0.656 N

b)

0.328 N

c)

1.312 N

d)

0.164 N

126.

What volume of 0.10 M Ba(NO3)2 solution is required to react completely with 100.0 mL of 0.50 M NaCl solution to form BaCl2(s)?

a)

250.0 mL

b)

200.0 mL

c)

500.0 mL

d)

100.0 mL

127.

The term colligative property refers to a property of a solution that depends only on the

a)

number of solute particles

b)

chemical nature of the solute

c)

solvent characteristics

d)

the number and type of solute particles

128.

Which statement correctly compares the Arrhenius and Brønsted–Lowry definitions of acids and bases?

a)

Arrhenius: Acids produce H+ ions in water, bases produce OH- ions.

b)

Arrhenius: Acids produce OH- ions in water, bases produce H+ ions.

c)

Brønsted–Lowry: Acids are proton donors, bases are proton acceptors.

d)

Brønsted–Lowry: Acids are proton acceptors, bases are proton donors.

129.

A conjugate acid–base pair in the Brønsted–Lowry model consists of two species that:

a)

differ by one proton (H+)

b)

have the same number of protons

c)

are both acids

d)

are both bases

e)

differ by one hydroxide ion (OH-)

130.

Which of the following are the balanced chemical equations showing HCl and H2SO4 behaving as Brønsted–Lowry acids in water?

a)

HCl + H2O → H3O+ + Cl-

H2SO4 + H2O → H3O+ + HSO4-

b)

HCl + H2O → H2O + Cl-

H2SO4 + H2O → H2O + SO42-

c)

HCl + H2O → HClO + H2

H2SO4 + H2O → H2SO3 + H2O2

131.

The strength of an acid is related to the position of its ionization equilibrium in that:

a)

the stronger the acid, the more the equilibrium lies to the right, indicating greater ionization.

b)

the stronger the acid, the more the equilibrium lies to the left, indicating less ionization.

c)

the strength of the acid does not affect the position of equilibrium.

d)

the weaker the acid, the more the equilibrium lies to the right, indicating greater ionization.

132.

Which of the following shows the correct dissociation (ionization) equations for HNO3 in water?

a)

HNO3 + H2O →

H3O+ + NO3-

b)

HNO3 + H2O →

H2O+ + HNO3-

c)

HNO3 + H2O →

H2O2 + HNO2-

d)

HNO3 + H2O →

H3O- + NO3+

133.

Water is considered an amphoteric substance because it can:

a)

act only as an acid

b)

act only as a base

c)

act as both an acid and a base

d)

neither donates nor accepts protons

134.

At 25 °C, what are the values of [H+], and [OH-] in pure water?

a)

They both equal 1.0×10−141.0 \times 10^{-14}

b)

They both equal 1.0×10−7M1.0 \times 10^{-7} M

c)

[OH-] = 1.0×10−7M1.0 \times 10^{-7} M

[H+] = 1.0×10−14M1.0 \times 10^{-14} M

d)

[H+] = 1.0×10−7M1.0 \times 10^{-7} M

[OH-] = 1.0×10−14M1.0 \times 10^{-14} M

135.

In an acidic solution, [H+] is ​ (a)   [OH-].

In a basic solution, [H+] is ​ (b)   [OH-].

Choose from the below words
the same as
greater than
less than
136.

The pH scale is defined as:

a)

pH = -log[H+]

b)

pH = [H+]

c)

pH = log[H+]

d)

pH = -[H+]

137.

Match the pH scales to the correct terms

a)

pH = 7

1.

Neutral

b)

pH < 7

2.

Acidic

c)

pH > 7

3.

Basic

138.

When the pH of a solution changes by one unit, by what factor does the hydrogen ion concentration change in the solution?

a)

2

b)

1

c)

10

d)

100

e)

log [H+]

139.

pOH is defined as:

a)

pOH = -log[OH-]

b)

pOH = -log[H+]

c)

pOH = log[OH-]

d)

pOH = log[H+]

140.

A buffered solution is important because it:

a)

resists changes in pH

b)

increases the rate of chemical reactions

c)

prevents the formation of acids and bases

d)

removes all impurities from solutions

141.

Write the conjugate acid of SO32- .

a)

HSO3-

b)

SO32-

c)

SO42-

d)

H2SO3

142.

Write the conjugate acid of F-.

a)

HF

b)

H2F

c)

FOH

d)

H2O

143.

Write the conjugate acid of CH3COO-.

a)

CH3COOH

b)

CH4

c)

HCOOH

d)

CH3OH

144.

Write the conjugate base of the following acid: H2SO4

a)

HSO4−

b)

HSO42-

c)

H3O+

d)

SO32-

145.

Write the conjugate base of the following acid: H2CO3

a)

HCO3−

b)

CO32-

c)

H2CO3−

d)

HCO2−

146.

For the following, calculate the indicated quantity: [H+]=4.01×10−3M;[H^+] = 4.01 \times 10^{-3} M; pH = ?

a)

pH = 2.397

b)

pH = 3.397

c)

pH = 1.397

d)

pH = 4.397

147.

For the following, calculate the indicated quantity: [OH−]=7.41×10−8M;[OH^{-}] = 7.41 \times 10^{-8} M; pOH = ?

a)

7.130

b)

6.870

c)

8.210

d)

5.920

148.

Calculate the pH and pOH values for the following solution: 0.00141 M HNO3

a)

pH = 2.85

pOH = 11.15

b)

pH = 3.85

pOH = 10.15

c)

pH = 1.85

pOH = 12.15

d)

pH = 4.85

pOH = 9.15

149.

Calculate the pH and pOH values for the following solution: 2.13×10−3MNaOH2.13 \times 10^{-3} M NaOH

a)

pOH = 2.67

pH = 11.33

b)

pOH = 11.33

pH = 2.67

c)

pOH = 3.67

pH = 10.33

d)

pOH = 1.33

pH = 12.67

150.

A 25.0-mL sample of 0.50 M HCl is titrated to the endpoint with 12.4 mL of NaOH. Calculate the concentration of the NaOH solution.

a)

1.01 M NaOH

b)

0.25 M NaOH

c)

0.51 M NaOH

d)

2.01 M NaOH

151.

pH and pOH for a given solution are related by which of the following equations at 25°C?

a)

pH + pOH = 7

b)

pH + pOH = 10

c)

pH + pOH = 14

d)

pH + pOH = 1

152.

Equal amounts of an acid and a base react completely, forming a salt and water. The pH of the resulting solution is checked, and rather than its being neutral, the solution is found to have a pH of 6.1. Using what you know about strong acids and bases, which of the following best categorizes the reacting acid and base?

a)

Strong acid and strong base

b)

Strong acid and weak base

c)

Weak acid and strong base

153.

The collision model for chemical reactions explains that what must "collide" for a reaction to occur?

a)

Atoms or molecules

b)

Electrons

c)

Neutrons

d)

Protons

154.

Activation energy is defined as:

a)

The minimum amount of energy required to start a chemical reaction.

b)

The energy released during a chemical reaction.

c)

The total energy of the reactants.

d)

The maximum amount of energy required to start a chemical reaction.

155.

A catalyst is a substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change. What do we call a biologic catalyst?

a)

Enzyme

b)

Hormone

c)

Vitamin

d)

Substrate

156.

A reaction "has reached a state of chemical equilibrium" when:

a)

the concentrations of reactants and products are equal

b)

the forward and reverse reaction rates are equal

c)

the reaction has stopped completely

d)

all reactants have been converted to products

157.

The equilibrium position in a chemical reaction refers to:

a)

The point where the concentrations of reactants and products remain constant.

b)

The point where all reactants are completely converted to products in a chemical reaction.

c)

The initial concentrations of reactants and products.

d)

The maximum rate of reaction.

158.

Homogeneous and heterogeneous equilibria differ in which of the following ways?

a)

The phases of reactants and products.

b)

The temperature of reactants and products.

c)

The enthalpy of formation of reactants and products.

d)

The concentration of reactants and products.

159.

Le Chatelier’s principle states that if a system at equilibrium is disturbed, the system will:

a)

Remain unchanged

b)

Shift to counteract the disturbance

c)

Speed up the reaction

d)

Stop the reaction completely

160.

Write the equilibrium constant expression for the following reaction: H2(g) + Br2(g) ⇌ 2HBr(g)

a)

Kc = [HBr]2[H2][Br2]\frac{[HBr]^2}{[H_2][Br_2]}

b)

Kc=[H2][Br2][HBr]2Kc=\frac{[H_2][Br_2]}{[HBr]^2}

c)

Kc = [H2]2[Br2]2[HBr]\frac{[H_2]^2[Br_2]^2}{[HBr]}

d)

Kc = [HBr][H2][Br2]\frac{[HBr]^{ }}{[H_2][Br_2]}

161.

Write the equilibrium constant expression for the following reaction: SO2Cl2(g) ⇌ SO2(g) + Cl2(g)

a)

[SO2][Cl2]

[SO2Cl2]

b)

[SO2Cl2]

[SO2][Cl2]

c)

[SO2Cl2][SO2][Cl2]

d)

1

[SO2][Cl2][SO2Cl2]

162.

Write the expression for Ksp for the following sparingly soluble salt: ZnS

a)

Ksp = [Zn2+][S2−]

b)

Ksp = [ZnS]

c)

Ksp = [Zn2+]/[S2−]

d)

Ksp = [Zn2+]2[S2−]2

163.

Write the expression for Ksp for the following sparingly soluble salt: HgCl2.

a)

Ksp=[Hg+2][Cl−]2K_{sp}=[Hg_{ }^{+2}][Cl^-]^2

b)

Ksp = [Hg2+][Cl-]

c)

Ksp = [Hg2+][Cl2]

d)

Ksp=[Hg+][Cl−]2K_{sp}=[Hg^+][Cl^-]^2

164.

Do all collisions between molecules result in the breaking of bonds and the formation of products? Why?

a)

Yes, all collisions result in bond breaking and product formation.

b)

No, only collisions with proper orientation and energy result in product formation.

c)

Yes, as long as molecules collide, products are always formed. from colliding reactants.

d)

No, collisions never result in bond breaking or product formation.

165.

How does the collision model account for the observation that higher concentrations and higher temperatures tend to make reactions occur faster?

a)

By increasing the number of effective collisions.

b)

By decreasing the activation energy required for the reaction.

c)

By increasing the activation energy required for the reaction.

d)

By reducing the energy of the reactant molecules.

166.

An increase in temperature for a reaction affects the number of collisions that possess energy greater than Ea by:

a)

Decreasing the number of such collisions

b)

Not affecting the number of such collisions

c)

Increasing the number of such collisions

d)

Eliminating all such collisions

167.

A catalyst speeds up a reaction by:

a)

Increasing the activation energy (Ea) of the reaction.

b)

Decreasing the activation energy (Ea) of the reaction.

c)

Changing the enthalpy of the reaction.

d)

Consuming itself in the reaction.

168.

Equilibrium is a dynamic state because:

a)

the forward and reverse reactions continue to occur at equal rates.

b)

no reactions are occurring at equilibrium.

c)

the concentrations of reactants and products are always changing.

d)

energy is not conserved at equilibrium.

169.

Which of the following changes will shift the equilibrium position of the reaction

2SO2(g) + O2(g) ⇌ 2SO3(g)

to the right?

a)

Adding SO2(g) to the system

b)

Removing SO2(g) from the system

c)

Using a very efficient catalyst

d)

Reducing the volume of the container drastically

170.

Dissolving a slightly soluble salt to form a saturated solution is an equilibrium process because:

a)

the rate of dissolution equals the rate of precipitation.

b)

all the salt dissolves completely.

c)

no ions are present in the solution.

d)

the solution becomes supersaturated.

171.

Silver chloride (AgCl) dissolves in water to give a solution containing 6.3 × 10−6 mol solute per liter at 10 °C.

Calculate Ksp for AgCl at this temperature.

a)

4.0 × 10−11

b)

1.2 × 10−8

c)

6.0 × 10−7

d)

2.2 × 10−10

172.

Which of the following changes will shift the equilibrium position in favor of additional products in a chemical system?

a)

Increasing the concentration of a reactant

b)

Removing one of the products

c)

Increasing the temperature for an endothermic reaction

d)

Increasing the temperature for an exothermic process

173.
What letter represents the activation energy?
a)
A
b)
B
c)
C
d)
D
174.

A + B ⟺ C + D + heatA\ +\ B\ \Longleftrightarrow\ C\ +\ D\ +\ heat
If energy is added to the reaction, the reaction will run ​ in the ​ (a)   direction.

Choose from the below words
reverse
forward
175.
For the reaction...
N2 (g) +  3 H2 (g) <−> 2 NH3 (g)

If the pressure in the system is increased, the reaction will __________________.
a)
 shift to the left
b)
shift to the right
c)
not shift
176.
An equilibrium constant with a large magnitude indicates…
a)
A very fast reaction
b)
More products at equilibrium
c)
More reactants at equilibrium
d)
nothing, without considering the stoichiometry of the reaction
177.

What is the activation energy for the forward reaction?

a)

80 kcal

b)

40 kcal

c)

20 kcal

d)

60 kcal

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