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AP 6 Thermochemistry Question Bank

Total questions: 206

Worksheet time: 3hrs 27mins

Name
Class
Date
1.

In an exothermic reaction, the sign of ΔH is ____ and the reaction feels ______

a)

(-), hot

b)

(+), hot

c)

(-), cold

d)

(+), cold

2.

Freezing is an _________________________ process

a)

Endothermic

b)

Exothermic

3.

Vaporizing is an _________________________ process

a)

Endothermic

b)

Exothermic

4.

Breaking bonds is _________________. Forming bonds is ______________________

a)

endothermic, exothermic

b)

exothermic, endothermic

5.

ΔHrxn = ΔH_________________ − ΔH_________________

a)

products, reactants

b)

reactants, products

6.

If a reaction is exothermic, then the bonds formed in the products are ___________________ than the reactant bonds.

a)

stronger/more stable

b)

weaker/less stable

7.

If a reaction is endothermic, then the bonds formed in the products are ___________________ than the reactant bonds.

a)

stronger/more stable

b)

weaker/less stable

8.

According to Hess' Law, if you double the coefficients of a reaction, ΔH will ________________________

a)

Stay the same

b)

Double

c)

Quadruple

d)

Halve

9.

According to Hess' Law, if you reverse a reaction, ΔH will ________________________

a)

Change signs

b)

Double

c)

Be inversed

d)

Halve

10.

_____ Hydrogen-Hydrogen single bond/s are broken and _____ Carbon-Hydrogen single bond/s formed

a)

6; 8

b)

2; 8

c)

2; 10

d)

6; 10

11.

_____ Carbon-Carbon single bond/s are formed

a)

6

b)

2

c)

3

d)

10

12.

Bond enthalpy is the amount of energy is takes to break __________ of a bond.

a)

one mole

b)

one gram

c)

the activation energy

d)

the binding energy

13.

What are the units of ΔHrxn?

a)

kJ

b)

J

c)

kJ/mol

d)

J/mol

14.

Endothermic

a)

absorb energy

b)

release energy

15.

Exothermic

a)

absorb energy

b)

release energy

16.

In the dissolution process, the increase in separation between ions in the solute is

a)

endothermic

b)

exothermic

17.

In a system, the energy lost by the reacting species must be ________________ the energy gained by the surrouding

a)

greater than

b)

less than

c)

equal to

18.

On a potential energy diagram for an endothermic reaction,

a)

the potential energy of the reactants is greater than the potential energy of the products

b)

the potential energy of the reactants is less than the potential energy of the products

c)

the potential energy of the reactants is equal to the potential energy of the products

19.

A temperature is placed in a calorimeter. A solid is dissolved in water in the calorimeter and the temperature of the thermometer increases. The dissolution reaction must be

a)

endothermic

b)

exothermic

20.

Using bond energies to solve for enthalpy, 

a)

ΔH\Delta H  =  Bond broken - bonds formed

b)

undefined =  Bond formed- bonds broken

c)

undefined =  products - reactants

d)

undefined =  reactants- products

21.

The enthalpy of formation for a lone element is

a)

a number

b)

zero

22.

For a classroom demonstration, a chemistry teacher puts samples of two different pure solid powders in a beaker. The teacher places the beaker on a small wooden board with a wet surface, then stirs the contents of the beaker. After a short time the students observe that the bottom of the beaker is frozen to the wood surface. The teacher asks the students to make a claim about the observation and to justify their claims. Which of the following is the best claim and justification based on the students’ observation?

a)

An exothermic chemical change occurred because heat flowed from the contents of the beaker to the room.

b)

An exothermic physical change occurred because heat flowed from the contents of the beaker and the water on the board to the room.

c)

An endothermic physical change occurred because the freezing of water is an endothermic process.

d)

An endothermic chemical change occurred because the temperature of the beaker and the water on the board decreased as heat was absorbed by the reaction.

23.

The dissolution of an ionic solute in a polar solvent can be imagined as occurring in three steps, as shown in the figure above. In step 1, the separation between ions in the solute is greatly increased, just as will occur when the solute dissolves in the polar solvent. In step 2, the polar solvent is expanded to make spaces that the ions will occupy. In the last step, the ions are inserted into the spaces in the polar solvent. Which of the following best describes the enthalpy change, ΔH, for each step?

a)

All three steps are exothermic.

b)

All three steps are endothermic.

c)

Steps 1 and 2 are exothermic, and the final step is endothermic.

d)

Steps 1 and 2 are endothermic, and the final step is exothermic.

24.

Which of the following phase changes involves the transfer of heat from the surroundings to the system?

a)

CH4(g)→CH4(l), because CH4 molecules in the gas phase must absorb energy in order to move closer together, thereby increasing the intermolecular attractions in the solid state.

b)

CO2(g)→CO2(s), because CO2 molecules in the gas phase must absorb energy in order to move closer together, thereby increasing the intermolecular attractions in the liquid state.

c)

H2O(l)→H2O(s), because H2O molecules in the liquid phase must absorb energy in order to create a crystalline structure with strong intermolecular attractions in the solid state.

d)

NH3(l)→NH3(g), because NH3 molecules in the liquid phase must absorb energy in order to overcome their intermolecular attractions and become free gas molecules.

25.

In the spring, blossoms on cherry trees can be damaged when temperatures fall below −2°C. When the forecast calls for air temperatures to be below −5°C for a few hours one night, a farmer sprays his blossoming cherry trees with water, claiming that the blossoms will be protected by the water as it freezes. Which of the following is a correct scientific justification for spraying water on the blossoms to protect them from temperatures below −2°C?

a)

Water on the blossoms will not freeze unless the air temperature falls significantly below −5°C

b)

Water is a good thermal conductor that transfers heat from the cold air to the blossoms, keeping the blossoms from going below −2°C

c)

The freezing of water is an endothermic process; thus, water that freezes on the blossoms absorbs heat from the atmosphere, which in turn keeps the blossoms above 0°C

d)

The freezing of water is an exothermic process; thus, water that freezes on the blossoms releases heat to keep the blossoms at or above −2°C

26.

K(s)+12Cl2(g)→KCl(s)             ΔH°=−437kJ/molrxnK(s)+\frac{1}{2}Cl_2(g)→KCl(s)\ \ \ \ \ \ \ \ \ \ \ \ \ ΔH°=-437kJ/mol_{rxn} The elements K and Cl react directly to form the compound KCl according to the equation above. Which of the values of ∆Ho for a process in the table is (are) less than zero (i.e., indicate(s) an exothermic process)?

a)

z only

b)

y and z only

c)

x, y, and z only

d)

w, x, y, and z

27.

When a 50.mL sample of C2H5OH is mixed with a 50.mL sample of H2O, the resulting mixture has a volume of 95mL, and the container is warm to the touch. Which of the following best describes these observations?

a)

A chemical reaction occurs, as evidenced by the volume of the resultant mixture being less than the total volume of the initial components.

b)

A chemical change occurs, as evidenced by the formation of new covalent bonds releasing more energy than is absorbed by the breaking of the existing covalent bonds.

c)

A physical change occurs, and the solvation process is exothermic.

d)

A physical change occurs, and the solvation process is endothermic.

28.

Which of the following best helps to explain why the value of ΔH° for the dissolving of CaF2 in water is positive?

a)

CaF2(s) is insoluble in water.

b)

CaF2(s) dissolves in water to form CaF2(aq) particles.

c)

Ca2+ ions have very strong ion-ion interactions with F- ions in the crystal lattice.

d)

Ca2+ ions have very strong ion-dipole interactions with water molecules in the solution.

29.

XY2→X+Y2XY_2→X+Y_2 The equation represents the decomposition of a compound XY2. The diagram shows two reaction profiles. Which of the following best describes the flow of heat when 1.0 mol of XY2 decomposes?

a)

50 kJ of heat is transferred to the surroundings.

b)

50 kJ of heat is transferred from the surroundings.

c)

100 kJ of heat is transferred to the surroundings.

d)

100 kJ of heat is transferred from the surroundings.

30.

A reaction profile shows....

a)

the relative difference in temperature of reactants and products

b)

the relative difference in amounts of reactants and products

c)

the relative difference in energy of reactants and products

d)

the relative difference in volume of reactants and products

31.

Which letter represents the transition state?

a)

A

b)

B

c)

C

d)

D

e)

E

32.

Which letter represents the potential energy of the products?

a)

A

b)

B

c)

C

d)

D

e)

E

33.

Which letter represents the potential energy of the reactants?

a)

A

b)

B

c)

C

d)

D

e)

E

34.

Which letter shows the enthalpy of the reaction?

a)

A

b)

B

c)

C

d)

D

e)

E

35.

What does 'X' represent in the energy profile shown here?

a)

Potential Energy

b)

Kinetic Energy

c)

Enthalpy

d)

Reaction progess (time)

36.

What does 'Z' represent in the energy profile shown here?

a)

Reactants Energy of Formation

b)

Kinetic Energy

c)

Activation Energy

d)

Products Energy of Formation

37.

What does 'W' represent in the energy profile diagram shown?

a)

Kinetic Energy

b)

Reactants Energy of Formation

c)

Products Energy of Formation

d)

Activation Energy

38.

What does 'Q' represent in the energy profile shown here?

a)

Potential Energy

b)

Reaction Progression (time)

c)

Reactants Energy of Formation

d)

Activation Energy

39.

Which of these physical changes is endothermic?

a)

Condensation

b)

Freezing

c)

Evaporation

d)

Deposition

40.
This chart shows what type of reaction?
a)
endothermic
b)
mesothermic
c)
exothermic
d)
hotothermic
41.
This chart shows what type of reaction?
a)
exothermic
b)
combustion
c)
endothermic
d)
mesothermic
42.
An energy profile for a reversible chemical reaction is shown. According to this profile, the activation energy of the reverse reaction is equal to
a)
D
b)
C - B
c)
A - B
d)
C- D
43.

In an energy profile diagram, what does the difference between the energy of the reactants and the energy of the products represent?

a)

Activation Energy

b)

Enthalpy Change

c)

Potential Energy

d)

Kinetic Energy

44.

A 1.0mol sample of He(g) at 25°C is mixed with a 1.0mol sample of Xe(g) at 50°C. Which of the following correctly predicts the changes in average kinetic energy and the average speed of the Xe(g) atoms that will occur as the mixture approaches thermal equilibrium?

a)

Average Kinetic Energy of Xe Atoms will INCREASE;

Average Speed of Xe Atoms will INCREASE

b)

Average Kinetic Energy of Xe Atoms will INCREASE;

Average Speed of Xe Atoms will decrease

c)

Average Kinetic Energy of Xe Atoms will decrease;

Average Speed of Xe Atoms will INCREASE

d)

Average Kinetic Energy of Xe Atoms will decrease;

Average Speed of Xe Atoms will decrease

45.

A 100g sample of a metal was heated to 100oC and then quickly transferred to an insulated container holding 100g of water at 22oC. The temperature of the water rose to reach a final temperature of 35oC. Which of the following can be concluded?

a)

The metal temperature changed more than the water temperature did; therefore the metal lost more thermal energy than the water gained.

b)

The metal temperature changed more than the water temperature did, but the metal lost the same amount of thermal energy as the water gained.

c)

The metal temperature changed more than the water temperature did; therefore the heat capacity of the metal must be greater than the heat capacity of the water.

d)

The final temperature is less than the average starting temperature of the metal and the water; therefore the total energy of the metal and water decreased.

46.

A 100 g sample of a metal was heated to 100oC and then quickly transferred to an insulated container holding 100 g of water at 22oC. The temperature of the water rose to reach a final temperature of 35oC. Which of the following can be concluded?

a)

The metal temperature changed more than the water temperature did; therefore the metal lost more thermal energy than the water gained.

b)

The metal temperature changed more than the water temperature did, but the metal lost the same amount of thermal energy as the water gained.

c)

The metal temperature changed more than the water temperature did; therefore the heat capacity of the metal must be greater than the heat capacity of the water.

d)

The final temperature is less than the average starting temperature of the metal and the water; therefore the total energy of the metal and water decreased.

47.

A 30.g sample of Al(s) is heated to 50°C and placed in a calorimeter containing 150g of water at 20°C. As the system approaches thermal equilibrium, energy is transferred

a)

from the Al(s) to the water and the temperature of the water decreases

b)

from the Al(s) to the water and the temperature of the Al(s) decreases

c)

from the water to the Al(s) and the temperature of the water increases

d)

from the water to the Al(s) and the temperature of the Al(s) increases

48.

A piece of Fe(s) at 25°C is placed into H2O(l) at 75°C in an insulated container. A student predicts that when thermal equilibrium is reached, the Fe atoms, being more massive than the H2O molecules, will have a higher average kinetic energy than the H2O molecules. Which of the following best explains why the student’s prediction is incorrect?

a)

At thermal equilibrium, the less massive H2O molecules would have a higher average kinetic energy than the Fe atoms because they are more free to move than are the Fe atoms.

b)

At thermal equilibrium, the collisions between the Fe atoms and the H2O molecules would cease because the average kinetic energies of their particles would have become the same.

c)

At thermal equilibrium, the movement of both the Fe atoms and the H2O molecules would cease; thus, the average kinetic energy of their particles would have to be the same.

d)

At thermal equilibrium, the average kinetic energy of the Fe atoms cannot be greater than that of the H2O molecules; the average kinetic energies must be the same according to the definition of thermal equilibrium.

49.

The graphs above show Maxwell-Boltzmann distributions for one-mole samples of Ar(g). Graph 1 shows the distribution of particle energies at 300K and graph 2 shows the distribution of particle energies at 600K. A student predicts that if the samples are combined in an insulated container and thermal equilibrium is attained, then the most probable particle energy will be between the most probable energy shown in graph 1 and the most probable energy shown in graph 2. Which of the following is the best justification for the student’s claim?

a)

When the samples are combined, the gas particles will collide with one another, with the net effect being that the speed of the lowest energy particles decreases while the speed of the highest energy particles increases, leaving the average speed of the particles in the original samples unchanged.

b)

When the samples are combined, the gas particles from each sample will collide with the gas particles from the other sample until every particle in the mixture has the same speed, which is between the average speed of the particles in the hotter sample and the average speed of the particles in the cooler sample.

c)

When the samples are combined, the gas particles collide with one another until every particle in the mixture has the same kinetic energy, which is between the average kinetic energy of the particles in the hotter sample and the average kinetic energy of the particles in the cooler sample.

d)

When the samples are combined, the gas particles will collide with one another, with the net effect being that energy will be transferred from the more energetic particles to the less energetic particles until a new distribution of energies is achieved at a temperature between 300K and 600K.

50.

A student adds 50.0g of liquid water at 25.0°C to an insulated container fitted with a temperature probe. The student then adds 10.0g of ice at 0.0°C, closes the container, and measures the temperature at different intervals. Part of the data is shown in the graph above. The student predicts that the temperature will continue to decrease then level out to a constant temperature. Which of the following best explains why the student’s prediction is correct?

a)

The H2O molecules initially in the ice and the molecules initially in the liquid will have the same average kinetic energy.

b)

The transfer of energy between the H2O molecules in the ice and liquid water stops once all the molecules are in the liquid phase.

c)

Once all of the H2O molecules are in the liquid phase, the individual molecular speeds either increase or decrease until all the particles have the same speed.

d)

Once all of the H2O molecules are in the liquid phase, collisions between them virtually stop as they reach an equilibrium distance from their neighboring molecules.

51.

When you heat up an object, the particles in the object move _________.

a)

faster

b)

the same speed

c)

slower

52.

The condition where two substances in physical contact with each other exchange no heat energy. They are said to be the same temperature.

a)

Thermal Energy

b)

Cooler to Warmer

c)

Energy Transfer

d)

Thermal Equilibrium

53.
heat always travels from _____ to ______ objects
a)
hot to cold
b)
cold to hot
c)
cold to cold
d)
hot to hot
54.

Two objects have different temperatures. Object A has a temperature of 42 degrees and Object B is 37 degrees. Which direction should the energy transfer between Objects A and B?

a)

From Object A to Object B

b)

From Object B to Object A

c)

Energy will not transfer

d)

Energy will transfer both directions

55.

What happens to the temperature if I add ice cubes to a glass of soda

a)

The heat transfers from the coke to the ice

b)

The heat transfers from the ice to the coke

56.

You leave your cold spoon in a hot bowl of oatmeal. What happens after an hour?

a)

The temperature of both the oatmeal and the spoon stay the same

b)

The spoon and the oatmeal become the same temperature

c)

Only the temperature of the oatmeal will change

d)

Only the temperature of the spoon will change

57.

Heating will continue until what is reached?

a)

Energizers

b)

Equilibrium

c)

Equality

d)

Energy

58.
Temperature is a measure of the...
a)
total energy in a substance
b)
total kinetic energy in a substance
c)
average potential energy in a substance
d)
average kinetic energy of molecules in a substance
59.
High temperature is connected to:
a)
Low kinetic energy.
b)
High kinetic energy.
c)
Zero net energy.
d)
Equal kinetic energies.
60.
Two objects, V and W, touching each other. The initial temperature of V and W are 90°C and 80°C respectively.
Which statement is correct when V and W are at thermal equilibrium?
a)
Net rate of heat flow between V and W is zero
b)
Rate of change in temperature of V is bigger then that of W
c)
Temperature of V is higher than W
d)
The quantity of heat in V is the same as in W
61.

A 1.0 g sample of a cashew was burned in a calorimeter containing 1000. g of water, and the temperature of the water changed from 20.0°C to 25.0°C. In another experiment, a 3.0 g sample of a marshmallow was burned in a calorimeter containing 2000. g of water, and the temperature of the water changed from 25.0°C to 30.0°C. Based on the data, which of the following can be concluded about the energy content for 1.0 g of each of the two substances? (The specific heat of water is 4.2 J/(g⋅°C).)

a)

The combustion of 1.0 g of cashew releases less energy than the combustion of 1.0 g of marshmallow.

b)

The combustion of 1.0 g of cashew releases the same amount of energy as the combustion of 1.0 g of marshmallow.

c)

The combustion of 1.0 g of cashew releases more energy than the combustion of 1.0 g of marshmallow.

d)

No comparison can be made because the two systems started with different masses of food, different masses of water, and different initial temperatures.

62.

NaOH(aq)+HCl(aq)→H2O(l)+NaCl(aq)NaOH(aq)+HCl(aq)→H_2O(l)+NaCl(aq) A student is trying to determine the heat of reaction for the acid-base neutralization reaction represented above. The student uses 0.50 M NaOH and 0.50 M HCl solutions. Which of the following situations, by itself, would most likely result in the LEAST error in the calculated value of the heat of reaction?

a)

The thermometer was incorrectly calibrated and read 0.5 Celsius degree too high during the procedure.

b)

The volume of the acid solution added to the calorimeter was actually 1.0 mL less than what was recorded.

c)

The calorimeter was poorly insulated, and some heat escaped to the atmosphere during the procedure.

d)

The actual molarity of the base solution was 0.53 M but was recorded as 0.50 M.

63.

Substance Specific Heat (J/(g⋅C∘)) Mass of Sample (g)

Fe(s) 0.46 1.0

Al(s) 0.91 1.0

Samples of Fe and Al are heated from 25∘C to 75∘C. Based on the information in the table above, the change in the average kinetic energy of the atoms in the samples is

a)

higher for the Fe sample, because with its lower specific heat it absorbs more energy

b)

higher for the Al sample, because with its higher specific heat it absorbs more energy

c)

the same for both metal samples, because the initial and final temperatures are the same

d)

the same for both metal samples, because the masses are equal

64.

Which of the following is a list of the minimum amount of data needed for determining the molar enthalpy of solution of KCl(s) in pure H2O(l) ? (Assume that the KCl(aq) has the same specific heat capacity as pure water and that the initial temperatures of the KCl(s) and the water are the same.)

a)

Mass of KCl(s), initial temperature of the water, and final temperature of the solution

b)

Mass of H2O, initial temperature of the water, and final temperature of the solution

c)

Mass of KCl(s), mass of H2O, initial temperature of the water, and final temperature of the solution

d)

Mass of KCl(s), mass of H2O, initial temperature of the water, final temperature of the solution, and atmospheric pressure

65.
For the formula:
 Q= m c ∆T
The  units for specific heat are:
a)
g /J C
b)
°C/g J
c)
kJ/g
d)
J/g°C
66.
What would be the effect on the particles if more heat is supplied to the system?
a)
They would slow down
b)
They would stop moving
c)
They would speed up
d)
There would be no effect
67.

If two objects have different temperatures when they come in contact, heat will flow from the warmer object to the cooler one UNTIL ____________

a)

one reaches a temperature of zero

b)

they both have an equal temperature

c)

one runs out of energy

68.
The total energy of all the particles in a substance is called: 
a)
temperature
b)
thermal energy
c)
degrees
d)
mass
69.

If the specific heat of water is 4.18 J/g∙°C, how much heat is required to increase the temperature of 1.2 kg of water from 23 °C to 39 °C? (show your work)

a)

-80,256 J

b)

80.256 J

c)

80,256 J

d)

-80.256 J

70.
What is the specific heat of an unknown substance if 100.0 g of it at 200.0 °C reaches an equilibrium temperature of 27.1 °C when it comes in contact with a calorimeter of water.  The water weighs 75. g and had an initial temperature of 20.00 °C?  (Specific heat of water is 4.18 J/g°C) (show your work)
a)
0.111 J/g°C
b)
1.29 J/g°C
c)
0.129 J/g°C
d)
22225.85 J
71.

For a skillet, used for cooking, do you want a high or low specific heat

a)

High, so that it will need more energy to heat up

b)

Low, so that it will change temperature quickly

72.
The SI unit of heat and energy is the __________.
a)
calorie
b)
heat
c)
joule
d)
watt
73.
A sample of iron receives 50.J of heat energy that raises the temperature of the iron by 25.0°C. If iron has a specific heat of 0.10 J/g°C, what is the mass of the iron sample?  (show your work)
a)
25g
b)
30g
c)
20g
d)
50g
74.
A reaction is performed in a beaker with a temperature probe recording the temperature changes of the reaction.  If the temperature began at 15.0 degrees Celsius and ended at 27.5 degrees Celsius. Is the reaction endothermic or exothermic?
a)
Exothermic
b)
Endothermic
75.
Calorimetry Question:
A piece of metal with a mass of 32.8 g is heated to 100.5°C and dropped into 138.2 g of water at 20.0°C.  The final temperature of the system is 30.2°C.  What is the specific heat capacity of the metal? (show your work)
a)
2.56 J/g°C
b)
0.391 J/g°C
c)
5.29 J/g°C
d)
3.50 J/g°C
76.
In an exothermic process, the surroundings are gaining energy.
a)
True
b)
False
77.

What does "q" mean?

a)

A measure of heat energy

b)

A change in heat energy

c)

A measure of kinetic energy

d)

A change in temperature

78.
What is a tool that measures the heat of chemical reactions?
a)
calorimeter
b)
 calorie
c)
Styrofoam™ cup
d)
none of these
79.
Calorimeters use insulating materials to:
a)
encourage affordable materials in science experiments in schools.
b)
preserve the heat of the environment outside the Styrofoam cup.
c)
prevent heat from escaping to the environment.
d)
keep the liquid inside the Styrofoam cup from overheating.
80.
A metal cube at temperature of 10°C immersed in a liquid at temperature of 70°C.
What is the temperature of the metal cube when thermal equilibrium is achieved between the cube and the liquid?
a)
Between 10°C and 70°C
b)
More than 70°C
c)
Less than 10°C
d)
Same as the room temperature
81.

The specific heat of water is 4.18 J/g°C.

If 980. J of energy is added to 6.20 g of water at 18.0 °C, what is the final temperature of the water?

a)

37.8 °C

b)

-19.8 °C

c)

19.8 °C

d)

55.8 °C

82.

What does "ΔT" mean?

a)

A change in health

b)

A change in heat

c)

A change in height

d)

A change in temperature

83.

In an experiment a student mixes a 50.0 mL sample of 0.100 M AgNO3(aq) with a 50.0 mL sample of 0.100 M NaCl(aq) at 20.0°C in a coffee-cup calorimeter. Which of the following is the enthalpy change of the precipitation reaction represented above if the final temperature of the mixture is 21.0°C? (Assume that the total mass of the mixture is 100. g and that the specific heat capacity of the mixture is 4.2 J/(g °C).) AgNO3+NaCl→AgCl(s)+NaNO3AgNO_3+NaCl→AgCl(s)+NaNO_3

a)

−84 kJ/molrxn

b)

−0.42 kJ/molrxn

c)

0.42 kJ/molrxn

d)

84 kJ/molrxn

84.

The heating curve for a sample of pure ethanol is provided above. The temperature was recorded as a 50.0 g sample of solid ethanol was heated at a constant rate. Which of the following explains why the slope of segment T is greater than the slope of segment R?

a)

The specific heat capacity of the gaseous ethanol is less than the specific heat capacity of liquid ethanol.

b)

The specific heat capacity of the gaseous ethanol is greater than the specific heat capacity of liquid ethanol.

c)

The heat of vaporization of ethanol is less than the heat of fusion of ethanol.

d)

The heat of vaporization of ethanol is greater than the heat of fusion of ethanol.

85.

In an insulated cup of negligible heat capacity, 50. g of water at 40.°C is mixed with 30. g of water at 20.°C. The final temperature of the mixture is closest to

a)

22°C

b)

27°C

c)

30.°C

d)

33°C

86.

A pen lying on a table possesses potential energy.

a)

TRUE

b)

FALSE

87.

If a person weighing 645 N climbs up a ladder to a height of 4.55 m, then what is the increase in gravitational potential energy?

a)

2,935 J

b)

645 J

c)

45 J

d)

28,790 J

88.

What is the change in gravitational potential energy of the apple if it falls the amount shown? Use g=9.81m/s² and round to the nearest hundredth.

a)

1.72 J

b)

0.70 J

c)

0.25 J

d)

0.18 J

89.

If a monkey of mass 8 kg climbs up a tree and the tree is 10 meters high, then the potential energy of the monkey at the top of the tree will be 784 J.

a)

TRUE

b)

FALSE

90.

A ball falls down 30 meters from the top of a building. If the ball weighed 1.2 kg, what is the gravitational potential energy lost by the ball? Estimate g as 9.81.

a)

456 J

b)

525 J

c)

300 J

d)

100 J

e)

353 J

91.

A stretched rubber band has potential energy.

a)

TRUE

b)

FALSE

92.

If the length of a steel wire increases by 0.2 cm on stretching it with a force of 20 N, then the work done in stretching the wire is (a)   .

93.

Sasha lifts a couch from the ground floor of her house to the attic. If the couch has a mass of 120 kg and is lifted 8.2 m, then what is the energy gained by the couch?

a)

9650 J

b)

11856 J

c)

1200 J

d)

1000 J

94.

If Sara has a mass of 60 kg and lives on the fifth floor of her apartment 18 m above the ground, then what is her gravitational potential energy?

a)

177 J

b)

589 J

c)

10595 J

d)

1000 J

95.

The ability to move an electric charge from one point to another is

a)

electrical kinetic energy

b)

potential energy

c)

electric potential energy

d)

kinetic energy

96.

What is the unit of measurement for potential energy?

a)

Watts

b)

Joules

c)

Newtons

d)

Meters

97.

The acceleration due to gravity (on Earth) is _____

a)

98 m/s2

b)

9.8 m/s2

c)

9.8 m/s

d)

0.98 m/s

98.
Roger Federer’s tennis ball has a mass of 0.3 kg. If he holds the ball above the ground at a height of 2.0 m to serve, what is its gravitational potential energy (GPE)?
a)
5.88 J
b)
0.6 J
c)
6.67 J
d)
0.15 J
99.
A leopard with a mass of 55.00 kg climbs 12.0 m up a tree. What is its GPE?
a)
6,468 J
b)
660 J
c)
4.58 J
d)
0.22 J
100.

Which picture is NOT an example of Chemical PE?

a)
b)
c)
d)
101.

The upper magnet is hung from the top of a box with a rubber band.

The bottom magnet is held in place at the bottom of a box.

What THREE potential energies are affecting this attraction in the picture?

a)

Gravitational, Electrical, Chemical

b)

Gravitational, Elastic, Chemical

c)

Magnetic, Elastic, Chemical

d)

Gravitational, Magnetic, Elastic

102.

How could you increase the gravitational potential energy of an object without changing its mass and gravity?

a)

Make the object larger

b)

Lower the object towards the ground

c)

Raise the object farther off the ground

d)

Allow the object to roll on the ground

103.

Elastic potential energy is energy stored in an object, such as a spring, due to the object being stretched or compressed. In which position is it most likely that the spring has lost its elastic potential energy?

a)

W

b)

Y

c)

X

d)

Z

104.

Which 2 items does gravitational pull depend on?

a)

Mass (weight)

b)

Speed

c)

Height / Distance from the ground

d)

Force

105.

True / False: The heavier an object is, the more potential energy the object has.

a)

True

b)

False

106.
Food has _______ potential energy that our bodies transform into kinetic energy.
a)
Elastic
b)
Nuclear 
c)
Gravitational
d)
Chemical
107.

At which position does the child on the swing have the most gravitational potential energy?

a)

W

b)

X

c)

Y

d)

Z

108.

Which object has elastic potential energy?

a)

a hanging leaf

b)

a burning coal

c)

a rolling ball

d)

a stretched bow

109.

What type of potential energy does this automobile battery contain?

a)

gravitational potential energy

b)

elastic potential energy

c)

chemical potential energy

d)

kinetic potential energy

110.

Explain what happens to the potential energy of a rock as it falls off the side of a cliff.

a)

The kinetic energy decreases as the potential energy increases.

b)

The kinetic energy decreases as the potential energy increases.

c)

The potential energy decreases as the kinetic energy increases.

d)

The potential energy decreases as the kinetic energy decreases.

111.

How much energy is required to melt 64 g of methane at 90 K? (The molar mass of methane is 16 g/mol.)

a)

0.24 kJ

b)

3.8 kJ

c)

33 kJ

d)

60. kJ

112.

A sample of CHCl3(s) was exposed to a constant source of heat for a period of time. The graph above shows the change in the temperature of the sample as heat is added. Which of the following best describes what occurs at the particle level that makes segment D longer than segment B?

a)

The specific heat capacity of the liquid is significantly higher than that of the solid, because the particles in the liquid state need to absorb more thermal energy to increase their average speed.

b)

The specific heat capacity of the solid is significantly higher than that of the gas, because the particles in the solid state need to absorb more thermal energy to increase their average speed.

c)

The enthalpy of fusion is greater than the enthalpy of vaporization, because separating molecules from their bound crystalline state requires more energy than separating molecules completely from the liquid state.

d)

The enthalpy of vaporization is greater than the enthalpy of fusion, because separating molecules completely from the liquid to form a gas requires more energy than separating molecules from their bound crystalline state to a liquid state.

113.

A 2.00mol sample of C2H5OH undergoes the phase transition illustrated in the diagram above. The molar enthalpy of vaporization, ΔHvap, of C2H5OH is +38.6kJ/mol. Which of the following best identifies the change in enthalpy in the phase transition shown in the diagram?

a)

+19.3kJ

b)

+77.2kJ

c)

−19.3kJ

d)

−77.2kJ

114.

The graph shows the temperature of a pure substance as it is heated at a constant rate in an open vessel at 1.0 atm pressure. The substance changes from the solid to the liquid to the gas phase. The substance is at its normal freezing point at time

a)

t1

b)

t2

c)

t3

d)

t4

e)

t5

115.

The graph shows the temperature of a pure substance as it is heated at a constant rate in an open vessel at 1.0 atm pressure. The substance changes from the solid to the liquid to the gas phase. Which of the following best describes what happens to the substance between t4 and t5?

a)

The molecules are leaving the liquid phase.

b)

The solid and liquid phases coexist in equilibrium.

c)

The vapor pressure of the substance is decreasing.

d)

The average intermolecular distance is decreasing.

e)

The temperature of the substance is increasing.

116.

The diagram above represents the melting of H2O(s). A 2.00mole sample of H2O(s) at 0°C melted, producing H2O(l) at 0°C. Based on the diagram, which of the following best describes the amount of heat required for this process and the changes that took place at the molecular level?

a)

3.01 kJ of heat was absorbed to decrease the average speed of the water molecules in the liquid, which decreases the distance between molecules.

b)

6.02 kJ of heat was absorbed to increase the number of hydrogen bonds between water molecules in the liquid compared to the solid.

c)

12.0 kJ of heat was absorbed to decrease the polarity of the water molecules, which increases the density of the liquid compared to the solid.

d)

12.0 kJ of heat was absorbed to overcome some of the hydrogen bonding forces holding the water molecules in fixed positions in the crystalline structure.

117.

A molecular solid coexists with its liquid phase at its melting point. The solid-liquid mixture is heated, but the temperature does not change while the solid is melting. The best explanation for this phenomenon is that the heat absorbed by the mixture

a)

is lost to the surroundings very quickly

b)

is used in overcoming the intermolecular attractions in the solid

c)

is used in breaking the bonds within the molecules of the solid

d)

causes the nonbonding electrons in the molecules to move to lower energy levels

e)

causes evaporation of the liquid, which has a cooling effect

118.

The cooling curve above shows how the temperature of a sample varies with time as the sample goes through phase changes. The sample starts as a gas, and heat is removed at a constant rate. At which time does the sample contain the most liquid?

a)

t1

b)

t2

c)

t3

d)

t4

e)

t5

119.

Between which points is the temperature of the substance remaining constant?

a)

A-B only

b)

A-B, C-D, and E-F

c)

B-C only

d)

B-C and D-E

120.

The melting point of the sample is

a)

-60 ºC

b)

20 ºC

c)

60 ºC

d)

100 ºC

121.

Which letter indicates where water is in both the solid and liquid phase at the same time?

a)

A

b)

B

c)

C

d)

D

e)

E

122.

In which region(s) does temperature increase?

a)

Region C only

b)

Regions B and D

c)

Regions A, C, and E

d)

Regions A and B

123.

In a heating curve, diagonal lines indicate temperature is __________, while flat lines indicate temperature is __________.

a)

Constant, changing

b)

Changing, constant

124.

Heat __________ depend on the amount of substance, while temperature __________.

a)

Does not, does

b)

Does, does not

125.

NaCl was heated starting at an initial temperature of 0°C. The heating curve is shown in the picture. What is the melting point of this substance?

a)

0°C

b)

801°C

c)

1000°C

d)

1465°C

126.

The reaction represented goes essentially to completion. The reaction takes place in a rigid, insulated vessel that is initially at 600 K

Which of the following statements about the bonds in the reactants and products is most accurate?

a)

The sum of the bond enthalpies of the bonds in the reactant is greater than the sum of the bond enthalpies of the bonds in the products.

b)

The sum of the bond enthalpies of the bonds in the reactant is less than the sum of the bond enthalpies of the bonds in the products.

c)

The length of the bond between carbon and oxygen in CH3OH is shorter than the length of the bond between carbon and oxygen in CO.

d)

All of the bonds in the reactant and products are polar.

127.

Which of the following statements must be true for the combustion of CH4(g), represented by the equation? CH4(g)+2 O2(g)→CO2(g)+2 H2O(g)            ΔH°<0CH_4(g)+2\ O_2(g)→CO_2(g)+2\ H_2O(g)\ \ \ \ \ \ \ \ \ \ \ \ ΔH°<0

a)

The bond dissociation energy between O atoms in O2 is greater than the sum of the bond dissociation energies of the two O−H bonds in H2O.

b)

More bonds are formed in the products than are broken in the reactants.

c)

The sum of the bond dissociation energies of the product molecules is EQUAL to the sum of the bond dissociation energies of the reactant molecules.

d)

The sum of the bond dissociation energies of the product molecules is GREATER than the sum of the bond dissociation energies of the reactant molecules.

128.

CO(g) + 2 H2(g) ⇄ CH3OH(g)       ΔH < 0

The synthesis of CH3OH(g) from CO(g) and H2(g) is represented by the equation above. The value of Kc for the reaction at 483 K is 14.5.

Which of the following statements is true about bond energies in this reaction?

a)

The energy absorbed as the bonds in the reactants are broken is greater than the energy released as the bonds in the product are formed.

b)

The energy released as the bonds in the reactants are broken is greater than the energy absorbed as the bonds in the product are formed.

c)

The energy absorbed as the bonds in the reactants are broken is less than the energy released as the bonds in the product are formed.

d)

The energy released as the bonds in the reactants are broken is less than the energy absorbed as the bonds in the product are formed.

129.

The oxidation of carbon monoxide can be represented by the chemical equation 2CO(g)+O2(g)→2CO2(g)2CO(g)+O_2(g)→2CO_2(g) . The table above provides the average bond enthalpies for different bond types. Based on the information in the table, which of the following mathematical expressions is correct for the estimated enthalpy change for the reaction?

a)

ΔHrxn=[2(1072 kJmol)+(498 kJmol)]−2(799 kJmol)ΔHrxn=[2(1072\ \frac{kJ}{mol})+(498\ \frac{kJ}{mol})]−2(799\ \frac{kJ}{mol})

b)

ΔHrxn=[2(1072 kJmol)+(498 kJmol)]−4(799 kJmol)ΔHrxn=[2(1072\ \frac{kJ}{mol})+(498\ \frac{kJ}{mol})]−4(799\ \frac{kJ}{mol})

c)

ΔHrxn=[2(799kJmol)+(142kJmol)]−4(360kJmol)ΔHrxn=[2(799\frac{kJ}{mol})+(142\frac{kJ}{mol})]−4(360\frac{kJ}{mol})

d)

ΔHrxn=[2(799kJmol)+(142kJmol)]−2(360kJmol)ΔHrxn=[2(799\frac{kJ}{mol})+(142\frac{kJ}{mol})]−2(360\frac{kJ}{mol})

130.

Shown above are the equation representing the decomposition of H2O2(l) and a table of bond enthalpies. On the basis of the information, which of the following is the enthalpy of decomposition of 2 mol of H2O2(l) ?

a)

−349kJ

b)

−203kJ

c)

203kJ

d)

349kJ

131.

An equation representing the dissociation of O2(g) and a table of bond enthalpies are shown above. Based on the information, which of the following is the enthalpy of dissociation for O2(g)?

a)

−641kJ/mol

b)

−495kJ/mol

c)

495kJ/mol

d)

641kJ/mol

132.

CH4(g) + Cl(g) → CH3(g) + HCl(g) ΔH° = -14 kJ/molrxn

NH3(g) + Cl(g) → NH2(g) + HCl(g) ΔH° = -36 kJ/molrxn

H2O(g) + Cl(g) → OH(g) + HCl(g) ΔH° = + 40 kJ/molrxn

Based on the data above, what can be concluded regarding the strength of the C-H, N-H, and O-H bonds in the molecules shown?

a)

The C-H bond is the strongest.

b)

The N-H bond is the strongest.

c)

The O-H bond is the strongest.

d)

Nothing can be concluded without knowing the strength of the H- Cl bond.

133.

A mixture of NO2(g) and N2O4(g) is placed in a glass tube and allowed to reach equilibrium at 70°C, as represented above.

Which of the following statements about ∆H° for the reaction is correct?

a)

ΔH° < 0 because energy is released when the N−N bond breaks.

b)

ΔH° < 0 because energy is required to break the N−N bond.

c)

ΔH° > 0 because energy is released when the N−N bond breaks.

d)

ΔH° > 0 because energy is required to break the N−N bond.

134.

Based on the bond energies shown in the table above, which of the following diagrams best represents the change in energy as the reaction represented below proceeds?

H2(g)+ Cl2(g)→2HCl(g)

a)

b)

c)

d)

135.

_____ Hydrogen-Hydrogen single bond/s are broken and _____ Carbon-Hydrogen single bond/s formed

a)

6; 8

b)

2; 8

c)

2; 10

d)

6; 10

136.

_____ Carbon-Carbon single bond/s are formed

a)

6

b)

2

c)

3

d)

10

137.

Bond enthalpy is the amount of energy is takes to break __________ of a bond.

a)

one mole

b)

one gram

c)

the activation energy

d)

the binding energy

138.

Using bond energies to solve for enthalpy, 

a)

ΔH\Delta H  =  Bond broken - bonds formed

b)

undefined =  Bond formed- bonds broken

c)

undefined =  products - reactants

d)

undefined =  reactants- products

139.
Some average bond enthalpies, in kJ/mol, are as follows:
H-H = 436
Cl-Cl = 242
H-Cl = 431
What is the enthalpy change for the decomposition of hydrogen chloride?
2 HCl --> H2 + Cl2
a)
-184 kJ
b)
+ 184 kJ
c)
- 247 kJ
d)
+ 247 kJ
140.

The average bond enthalpies for O-O and O=O are 146kJ/mol and 496kJ/mol respectively. Calculate the enthalpy change for the reaction shown.

a)

-102kJ

b)

+102kJ

c)

+146kJ

d)

-248 kJ

141.

Breaking bonds is _________________. Forming bonds is ______________________

a)

endothermic, exothermic

b)

exothermic, endothermic

142.
Using the table of average bond energies below the image, the ΔH for the reaction in kJ is.....
a)
+ 160 kJ
b)
-63 kJ
c)
- 160 kJ
d)
-217 kJ
143.

Using the reaction & table above, determine the change in energy and whether or not the reaction is endothermic or exothermic:

a)

-2054 kJ endothermic

b)

2023 kJ endothermic

c)

-2023 kJ exothermic

d)

2023 kJ exothermic

144.

Use the bond energy data below to find the  ΔH  for 

2H2O(g) ⟶  2H2(g)  +  O2(g)

a)

  +480 kJ

b)

  -480 kJ

c)

  -440 kJ

d)

  + 440 kJ

145.

An equation representing the dissociation of O2(g) into two oxygen atoms, and a table of bond enthalpies are shown above. Based on the information, which of the following is the enthalpy of dissociation for O2(g)?

a)

−641kJ/mol

b)

−495kJ/mol

c)

495kJ/mol

d)

641kJ/mol

146.

3 C2H2(g) → C6H6(g)

What is the standard enthalphy change ΔHo, for the reaction represented above?

(ΔHof of C2H2(g) is 230 kJ mol-1;

(ΔHof of C6H6(g) is 83 kJ mol-1;)

a)

-607 kJ

b)

-147 kJ

c)

-19 kJ

d)

+19 kJ

e)

+773 kJ

147.

A skier starts from rest at the top of a 45.0 m hill and coasts down a 30° slope into a valley, then continues up to the top of a 40.0 m hill. Assuming no friction, how fast will the skier be moving at the top of the 40.0 m hill?

a)

4.6 m/s

b)

9.9 m/s

c)

2.5 m/s

d)

3.6 m/s

148.

The Law of Conservation of Energy states that the total energy of a system is always conserved. Is this statement:

a)

TRUE

b)

FALSE

149.

A 0.5 kg ball is at the top of a ramp which is 8 meters high. How much kinetic energy does the ball have at the bottom of the ramp?

a)

40 J

b)

4 J

c)

10 J

d)

400 J

150.

If an object is dropped from a certain height, does the total energy of that object keep decreasing until it becomes zero once it strikes the ground?

a)

TRUE

b)

FALSE

151.

How fast will the skier be moving on the valley floor between two hills of heights 45.0 m and 40.0 m, assuming no friction and no effort into the motion?

a)

24.8 m/s

b)

12.4 m/s

c)

9.9 m/s

d)

4.6 m/s

152.

A roller coaster begins at rest 120 m above the ground, as shown in the diagram. Assuming no friction from the wheels and air, and that no energy is lost to heat, sound, and so on, the radius of the loop is 40 m. Is the speed of the roller coaster at points H close to zero?

a)

TRUE

b)

FALSE

153.

If the velocity of an airplane is doubled, then its kinetic energy will increase by a factor of ____.

a)

8

b)

It will remain the same

c)

4

d)

2

154.

If a 2.00 kg ball is thrown straight upward and its maximum height is 25.51 m, approximately how much potential energy does the ball possess when it reaches its maximum height?

a)

200 J

b)

500 J

c)

100 J

d)

400 J

155.

A basketball player was falling to the ground after dunking the winning basket. At the end of her fall, she was falling at 4 m/s. If she was 60 kg, how much potential energy did she have at the top of her jump?

a)

240 J

b)

480 J

c)

960 J

d)

1200 J

156.

What is the law of conservation of energy?

a)

Energy can be transferred or transformed, but the total amount of energy remains constant.

b)

Energy can be created and destroyed in chemical reactions.

c)

Energy is always lost as heat in any transformation.

d)

Energy can be stored indefinitely without any losses.

157.

What energy transformations are occurring in this picture?

a)

gravitational potential --> sound --> thermal

b)

chemical potential --> mechanical --> elastic potential --> gravitational

c)

electrical --> sound --> thermal --> radiant

d)

chemical potential --> radiant --> thermal

158.

The sum of kinetic and potential energy in an object that is used to do work.

a)

Elastic Energy

b)

Sound Energy

c)

Magnetic Energy

d)

Mechanical Energy

159.
What is the transformation shown in the picture?
a)
chemical to electrical to radiant
b)
electrical to chemical to radiant
c)
radiant to chemical to electrical
d)
chemical to radiant to electrical
160.
Wind turbines and windmills convert (or change) what energy to another?
a)
Kinetic to electric
b)
Gravitational potential to electric
c)
Mechanical to potential
d)
Thermal to sound
161.

Which of the following best explains why an explosive charge in a cannon can cause a cannonball to fly out of the cannon with a high velocity?

a)

because chemical potential energy can convert to gravitational potential energy

b)

because chemical potential energy can convert to kinetic energy

c)

because chemical potential energy can convert to sound energy

d)

because chemical potential energy can convert to electric potential energy

162.

As an object falls in a gravitational field, its speed increases. This is an example of potential energy transforming into what?

a)

A.chemical potential energy

b)

electrical potential energy

c)

elastic potential energy

d)

kinetic energy

163.

Photosynthesis goes through two energy transformations. What are the transformations?

a)

radiant

chemical

b)

chemical

radiant

c)

electrical

nuclear

d)

radiant

sound

164.

A hot air ballon goes through three energy transformations. What are the transformations?

a)

chemical

thermal

mechanical

b)

mechanical

thermal

chemical

c)

chemical

mechanical

thermal

d)

thermal

chemical

mechanical

165.

What energy transformation occurs in the chameleon as it quickly launches its tongue to catch an insect?

a)

Radiant --> Chemical

b)

Thermal --> Chemical

c)

Chemical --> Mechanical

d)

Mechanical --> Chemical

166.
A television works by transforming electrical energy into what other kinds of energy?
a)
Light, thermal, and sound
b)
Thermal, chemical, and solar
c)
Sound, mechanical, and electrical
d)
Light, chemical, and mechanical
167.

Based on the information for two different reactions given above, which of the following gives the quantities needed to calculate the enthalpy change for the reaction represented by the overall equation below?

2NO(g)+O2(g)→N2O4(g)

a)

ΔH1+ΔH2

b)

ΔH1+(−ΔH2)

c)

(−ΔH1)+ΔH2

d)

ΔH1+(2×ΔH2)

168.

A → X

The enthalpy change for the reaction represented above is ΔHT. This reaction can be broken down into a series of steps as shown in the diagram:

A relationship that must exist among the various enthalpy changes is...

a)

ΔHT – ΔH1 – ΔH2 – ΔH3 = 0

b)

ΔHT + ΔH1 + ΔH2 + ΔH3 = 0

c)

ΔH3 – (ΔH1 + ΔH2) = ΔHT

d)

ΔH2 – (ΔH3 + ΔH1) = ΔHT

e)

ΔHT + ΔH2 = ΔH1 + ΔH3

169.

What is the enthalpy change for the sublimation of iodine, represented below?

I2(s) → I2(g)

a)

15 kJ/molrxn

b)

21 kJ/molrxn

c)

31 kJ/molrxn

d)

42 kJ/molrxn

e)

62 kJ/molrxn

170.

H2(g) + Cl2(g) ⇄ 2 HCl(g)      Kp = 2 × 1030 at 298 K 

HCl(g) can be synthesized from H2(g) and Cl2(g) as represented above. A student studying the kinetics of the reaction proposes the pictured mechanism.

What is the value of the enthalpy change per mole of HCl(g) produced?

a)

−93 kJ

b)

−121 kJ

c)

−186 kJ

d)

−242 kJ

171.

Based on the chemical equations and their associated enthalpy changes shown above, which of the following identifies the quantities needed to calculate ΔH°f, the standard enthalpy of formation of H2O(l), in kJ/mol?

a)

2(−ΔH°1) + (−ΔH°2) + 2(ΔH°3) + 2(ΔH°4)

b)

(−ΔH°1) + 1/2(−ΔH°2) + (ΔH°3) + (ΔH°4)

c)

(−ΔH°1) + 1/2(ΔH°2) + (ΔH°3)

d)

(ΔH°1) + 1/2(ΔH°2) + (ΔH°3) + (ΔH°4)

172.

CS2(l)+2H2O(l)→CO2(g)+2H2S(g)       ΔH°rxn=?CS_2(l)+2H_2O(l)→CO_2(g)+2H_2S(g)\ \ \ \ \ \ \ ΔH°rxn=?

Which of the following combinations represents the individual reactions and the quantities needed to determine ΔH° for the overall reaction represented by the chemical equation above?

a)

b)

c)

d)

173.

K(s)+ ½ Cl2(g) → KCl(s)                     ΔH° = -437kJ/molrxn

The elements K and Cl react directly to form the compound KCl according to the equation above. Refer to the information above and the table to answer the questions that follow.

Cl2(g) + 2e- → 2Cl-(g)

Which of the following expressions is equivalent to ΔHo for the reaction represented above?

a)

x + y

b)

x - y

c)

x + 2y

d)

x2−y

174.

Using the equations below:

C(s) + O2(g) → CO2(g) ∆H = –390 kJ

Mn(s) + O2(g) → MnO2(s) ∆H = –520 kJ

what is ∆H (in kJ) for the following reaction?

MnO2(s) + C(s) → Mn(s) + CO2(g)

a)

910

b)

130

c)

-130

d)

-910

175.

Using the equations below

Cu(s) + 1/2O2(g) → CuO(s) ∆H = –156 kJ

2Cu(s) + O2(g) → Cu2O(s) ∆H = –170 kJ

what is the value of ∆H (in kJ) for the following reaction?

2CuO(s) → Cu2O(s) + 1/2O2(g)

a)

142

b)

15

c)

-15

d)

-142

176.

Consider the following equations.

Mg(s) + O2(g) → MgO(s) ∆H = –602 kJ

H2(g) + O2(g) → H2O(g) ∆H = –242 kJ

What is the ∆H value (in kJ) for the following reaction?

MgO(s) + H2(g) → Mg(s) + H2O(g)

a)

-844

b)

-360

c)

+360

d)

+844

177.

The following equations show the oxidation of carbon and carbon monoxide to carbon dioxide.

C(s) +O2(g) → CO2(g) ΔH = –x kJ mol–1

CO(g) + O2(g) → CO2(g) ΔH = –y kJ mol–1

What is the enthalpy change, in kJ mol–1, for the oxidation of carbon to carbon monoxide?

C(s) + O2(g) → CO(g)

a)

x + y

b)

-x - y

c)

y - x

d)

x - y

178.

Which of the following statements are true for the reaction:

SO2(g) + 1/2O2(g) ↔ SO3(g)

ΔH = –92 kJ mol-1

Where ↔ indicates that the reaction can proceed in the forward and the reverse direction.

a)

The forward and reverse reaction both produce 92 kJ of energy.

b)

Oxidizing 2 moles of SO2 would produce twice as much energy.

c)

The reverse reaction has an enthalpy of +92 kJ mol-1.

d)

Collecting the SO3 produced in the liquid state would not change the measured enthalpy.

179.

The standard enthalpy change of formation values of two oxides of phosphorus are:

P4(s) + 3O2(g) → P4O6(s) ΔHf = –1600 kJ mol–1

P4(s) + 5O2(g) → P4O10(s) ΔHf = –3000 kJ mol–1

What is the enthalpy change, in kJ mol–1, for the reaction below?

P4O6(s) + 2O2(g) → P4O10(s)

a)

+4600

b)

+1400

c)

–1400

d)

–4600

180.

The enthalpies of combustion of C(s), H2(g) and C4H9OH(l) (in kJmol-1) are as follows   C(s) + O2(g)   --->  CO2(g)   ∆H=a H2(g) + ½O2(g)   --->   H2O(l)  ∆H=b C4H9OH(l) + 6O2(g)   --->   4CO2(g) + 5H2O(l)  ∆H=c

What is the enthalpy change for the reaction shown below?   4C(g) + 5H2(l) + ½O2(g)   --->   C4H9OH(l)

a)
c – 4a – 5b
b)
2a + 10b - c
c)
4a + 5b - c
d)
2a + 5b + c
181.

If N2 (g) + 2O2 (g) ⟶\longrightarrow   2NO2(g) has a ΔHrxn \Delta H_{rxn\ }  = 68, then what is the ΔHrxn\Delta H_{rxn}  if you reverse the reaction?

a)

86 kJ

b)

- 86 kJ

c)

68 kJ

d)

-68 kJ

182.

How much energy is required to turn 1 mole of N2O4(g) into 2 moles N & 4 moles O(g)?

a)

1933 kJ

b)

-1875 kJ

c)

- 1933 kJ

d)

1875 kJ

183.

Using the equations below:

C(s) + O2(g) → CO2(g) ∆H = –400 kJ

Mn(s) + O2(g) → MnO2(s) ∆H = –500 kJ

what is ∆H (in kJ) for the following reaction?

MnO2(s) + C(s) → Mn(s) + CO2(g) ∆H = ??? kJ

a)

-900 kJ

b)

900 kJ

c)

–100 kJ

d)

100 kJ

184.

Using the equations below

Cu(s) + 1/2O2(g) → CuO(s) ∆H = –156 kJ

2Cu(s) + O2(g) → Cu2O(s) ∆H = –170 kJ

what is the value of ∆H (in kJ) for the following reaction?

CuO(s) → Cu(s) + 1/2O2(g) ∆H = ???? kJ

a)

156 kJ

b)

170 kJ

c)

326 kJ

d)

14 kJ

185.

The standard enthalpy change of formation values of two oxides of phosphorus are:


P4(s) + 3O2(g) → P4O6(s) ΔHf = –1600 kJ mol–1

P4(s) + 5O2(g) → P4O10(s) ΔHf = –3000 kJ mol–1


What is the enthalpy change, in kJ mol–1, for the reaction below?


P4O6(s) + 2O2(g) → P4O10(s)

a)

+4600

b)

+1400

c)

–1400

d)

–4600

186.

The following equations show the oxidation of carbon and carbon monoxide to carbon dioxide.


C(s) +O2(g) → CO2(g) ΔH = –x kJ mol–1

CO(g) + O2(g) → CO2(g) ΔH = –y kJ mol–1


What is the enthalpy change, in kJ mol–1, for the oxidation of carbon to carbon monoxide?


C(s) + O2(g) → CO(g)

a)

x + y

b)

-x - y

c)

y - x

d)

x - y

187.

From the following data,

H2 (g) + Cl2 (g) → 2HCl(g) ΔH° = -185 kJ,

2H2 (g) + O2 (g) →2H2O(g) ΔH° = -484 kJ

calculate ΔH° for the following reaction.

4HCl(g) + O2 (g) →2Cl2 (g) + 2H2O(g)

a)

299 kJ

b)

-114 kJ

c)

-299 kJ

d)

114 kJ

188.

Using the equation below:

C(s) + O2(g) → CO2(g) ∆H = –400 kJ

what is ∆H (in kJ) for the following reaction?

CO2(g) → C(s) + O2(g) ∆H = ?? kJ

a)

+400 kJ

b)

–400 kJ

c)

1400 kJ\frac{1}{400}\ kJ

d)

–800 kJ

189.

Using the equations below:

Mn(s) + O2(g) → MnO2(s) ∆H = –500 kJ

what is ∆H (in kJ) for the following reaction?

MnO2(s) → Mn(s) + O2(g) ∆H = ??? kJ

a)

+500 kJ

b)

–500 kJ

c)

1500 kJ\frac{1}{500}\ kJ

d)

−1500 kJ-\frac{1}{500}\ kJ

190.

Using the equations below:

Mn(s) + O2(g) → MnO2(s) ∆H = –500 kJ

what is ∆H (in kJ) for the following reaction?

3 MnO2(s) → 3 Mn(s) + 3 O2(g) ∆H = ??? kJ

a)

+ 500 kJ

b)

–500 kJ

c)

–1500 kJ

d)

+ 1500 kJ

191.

Using the equation below:

C(s) + O2(g) → CO2(g) ∆H = –400 kJ

what is ∆H (in kJ) for the following reaction?

4 CO2(g) → 4 C(s) + 4 O2(g) ∆H = ?? kJ

a)

+400 kJ

b)

–400 kJ

c)

–1600 kJ

d)

+ 1600 kJ

192.

Consider the following equation:

2 Mg(s) + O2(g) → 2 MgO(s) ∆H = –1200 kJ

What is the ∆H value (in kJ) for the following reaction?

4 Mg(s) + 2 O2(g) → 4 MgO(s) ∆H = ??? kJ

a)

-2400 kJ

b)

+2400 kJ

c)

-4800 kJ

d)

4800 kJ

193.

Using the equations below:

C(s) + O2(g) → CO2(g) ∆H = –400 kJ

Mn(s) + O2(g) → MnO2(s) ∆H = –500 kJ

what is ∆H (in kJ) for the following reaction?

MnO2(s) + C(s) → Mn(s) + CO2(g) ∆H = ??? kJ

a)

-900 kJ

b)

900 kJ

c)

–100 kJ

d)

100 kJ

194.

Using the equations below

Cu(s) + 1/2O2(g) → CuO(s) ∆H = –156 kJ

2Cu(s) + O2(g) → Cu2O(s) ∆H = –170 kJ

what is the value of ∆H (in kJ) for the following reaction?

CuO(s) → Cu(s) + 1/2O2(g) ∆H = ???? kJ

a)

156 kJ

b)

170 kJ

c)

326 kJ

d)

14 kJ

195.

The standard enthalpy change of formation values of two oxides of phosphorus are:


P4(s) + 3O2(g) → P4O6(s) ΔHf = –1600 kJ mol–1

P4(s) + 5O2(g) → P4O10(s) ΔHf = –3000 kJ mol–1


What is the enthalpy change, in kJ mol–1, for the reaction below?


P4O6(s) + 2O2(g) → P4O10(s)

a)

+4600

b)

+1400

c)

–1400

d)

–4600

196.

The following equations show the oxidation of carbon and carbon monoxide to carbon dioxide.


C(s) +O2(g) → CO2(g) ΔH = –x kJ mol–1

CO(g) + O2(g) → CO2(g) ΔH = –y kJ mol–1


What is the enthalpy change, in kJ mol–1, for the oxidation of carbon to carbon monoxide?


C(s) + O2(g) → CO(g)

a)

x + y

b)

-x - y

c)

y - x

d)

x - y

197.

From the following data,

H2 (g) + Cl2 (g) → 2HCl(g) ΔH° = -185 kJ,

2H2 (g) + O2 (g) →2H2O(g) ΔH° = -484 kJ

calculate ΔH° for the following reaction.

4HCl(g) + O2 (g) →2Cl2 (g) + 2H2O(g)

a)

299 kJ

b)

-114 kJ

c)

-299 kJ

d)

114 kJ

198.

Using the equation below:

C(s) + O2(g) → CO2(g) ∆H = –400 kJ

what is ∆H (in kJ) for the following reaction?

CO2(g) → C(s) + O2(g) ∆H = ?? kJ

a)

+400 kJ

b)

–400 kJ

c)

1400 kJ\frac{1}{400}\ kJ

d)

–800 kJ

199.

Using the equations below:

Mn(s) + O2(g) → MnO2(s) ∆H = –500 kJ

what is ∆H (in kJ) for the following reaction?

MnO2(s) → Mn(s) + O2(g) ∆H = ??? kJ

a)

+500 kJ

b)

–500 kJ

c)

1500 kJ\frac{1}{500}\ kJ

d)

−1500 kJ-\frac{1}{500}\ kJ

200.

Using the equations below:

Mn(s) + O2(g) → MnO2(s) ∆H = –500 kJ

what is ∆H (in kJ) for the following reaction?

3 MnO2(s) → 3 Mn(s) + 3 O2(g) ∆H = ??? kJ

a)

+ 500 kJ

b)

–500 kJ

c)

–1500 kJ

d)

+ 1500 kJ

201.

Using the equation below:

C(s) + O2(g) → CO2(g) ∆H = –400 kJ

what is ∆H (in kJ) for the following reaction?

4 CO2(g) → 4 C(s) + 4 O2(g) ∆H = ?? kJ

a)

+400 kJ

b)

–400 kJ

c)

–1600 kJ

d)

+ 1600 kJ

202.

Using the equations below:

Mn(s) + O2(g) → MnO2(s) ∆H = –500 kJ

what is ∆H (in kJ) for the following reaction?

3 Mn(s) + 3 O2(g) → 3 MnO2(s) ∆H = ??? kJ

a)

+ 500 kJ

b)

–500 kJ

c)

–1500 kJ

d)

+ 1500 kJ

203.

Consider the following equation:

2 Mg(s) + O2(g) → 2 MgO(s) ∆H = –1200 kJ

What is the ∆H value (in kJ) for the following reaction?

4 Mg(s) + 2 O2(g) → 4 MgO(s) ∆H = ??? kJ

a)

-2400 kJ

b)

+2400 kJ

c)

-4800 kJ

d)

4800 kJ

204.

Consider the following equation:

2 Mg(s) + O2(g) → 2 MgO(s) ∆H = –1200 kJ

What is the ∆H value (in kJ) for the following reaction?

2 MgO(s) → 2 Mg(s) + O2(g) ∆H = ??? kJ

a)

-1200 kJ

b)

+ 1200 kJ

c)

11200 kJ\frac{1}{1200}\ kJ

d)

−11200kJ-\frac{1}{1200}kJ

205.

Consider the following equation.

2 H2(g) + O2(g) → 2 H2O(g) ∆H = –500 kJ

What is the ∆H value (in kJ) for the following reaction?

4 H2(g) + 2 O2(g) → 4 H2O(g) ∆H = ??? kJ

a)

–1000 kJ

b)

+1000 kJ

c)

+1500kJ+\frac{1}{500}kJ

d)

−1500kJ-\frac{1}{500}kJ

206.

Using the equations below:

C(s) + O2(g) → CO2(g) ∆H = –400 kJ

Mn(s) + O2(g) → MnO2(s) ∆H = –500 kJ

what is ∆H (in kJ) for the following reaction?

Mn(s) + CO2(g) → MnO2(s) + C(s) ∆H = ??? kJ

a)

-900 kJ

b)

900 kJ

c)

–100 kJ

d)

100 kJ

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Week 18 FINAL - Chemistry

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Week 18 FINAL - Chemistry

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10th - 12th Grade