AP Chemistry Packet



Chapter 5 Targets - Thermochemistry

| Target |Section in Book |Book problems |

| | |associated with |

| | |this target |

|1. I can give examples of different forms of energy. I can also state common units of |5.1 & 5.2 |13, 14 & 16 |

|energy and convert between these units. | | |

|2. I can define the first law of thermodynamics and write the associated equation. |5.2 |25 & 26 |

|3. I can describe how the change in internal energy of a system is related to the |5.2 |27 & 28 |

|changes of heat and work between the system and its surroundings. | | |

|4. I can define enthalpy and relate the enthalpy change in a process to the heat |5.2 & 5.3 |33 & 34 |

|added (endothermic) to or lost (exothermic) by the system during the process. | | |

|5. I can define the term state function and be provide examples of state functions. |5.2 |31 & 32 |

|6. I can sketch an exothermic and endothermic energy diagram and label the |5.2 |39 & 40 |

|various parts. | | |

|7. I can make stoichiometric calculations based upon a thermochemical equation. |5.4 |43, 45 & 47 |

|8. I can differentiate between the terms heat capacity and specific heat. |5.5 |49 & 50 |

|9. I can solve problems using q = m•Cp•∆T; calculate any one of the quantities |5.5 |53, 54 & 56 |

|given the other three. | | |

|10. I can describe the difference between a coffe cup calorimeter and a bomb |5.5 |57-60 |

|calorimeter. I will also be able to solve simple bomb calorimetry problems. | | |

|11. I can state Hess’s law and use Hess’s Law to solve problems. |5.6 |61, 63, 65 & 66 |

|12. I can define the terms standard state and standard heat of formation. I will also |5.7 |67, 68 & 70 |

|be to write a chemical reaction associated with the standard heat of formation. | | |

|13. I can calculate the enthalpy change in a reaction when given the standard |5.7 |73, 75, 78 & 80 |

|enthalpies of formations of each reactant and product. | | |

Chapter 5 - Thermochemistry

Thermochemistry is the study of the relationships between chemical reactions and energy changes.

|Target #1: I can give examples of different forms of energy. I can also state common units of energy and convert between these units. |

Energy – the capacity to do work or to transfer heat

Kinetic Energy - energy of motion; KE = 1/2 mv2; since all atoms have mass and are in motion, all atoms have a certain amount of kinetic energy; the energy that a substance has because of its temperature is related to its KE

Potential Energy - the energy that an object possess because of its composition or its position with respect to another object

Electrostatic potential energy – a form of potential energy due to the interactions between charged particles; this energy increases with increasing charge and decreases with increasing distance between the charged particles

Internal energy of a system (E) - sum of all KE and PE in a system; the difference in changes of energy from before the change to after the change; ∆E = Efinal – Einitial;

Relevance of the signs of ∆E:

positive ∆E = system has gained energy from the surroundings

negative ∆E = system has lost energy to its surroundings

• It is nearly impossible to determine the energy of a system, but it is relatively easy to calculate the changes of energy in a system!

Common Units of Energy:

A- joule; 1 J = 1 kg-m2/s2

B- calorie - used in biochemistry; 1 cal = 4.184 J

C- kilocalorie (Calorie); 1000 cal = 1 kilocalorie = 1 Calorie

Other terms related to energy changes . . .

System: portion of the universe that we are studying with respect to thermochemical changes

Surroundings: everything that lies outside of the system

example: If we were studying the combustion of gasoline inside of a cylinder, the gasoline and the oxygen would constitute the system and the cylinder and piston (and everything outside of that) would constitute the surroundings

Closed System: can exchange energy but not matter with the surroundings

Open System: can exchange both energy and matter with the surroundings

Heat: energy transferred from a hotter object to a colder object

Energy: capacity of an object to do work or to transfer heat (extensive property)

|Target #2: I can define the first law of thermodynamics and write the associated equation. |

First Law of Thermodynamics - energy is conserved; any energy lost by a system must

be gained by its surroundings (and visa-versa)

Equation representing the First law of Thermodynamics:

∆E = q + w, where . . .

E = change in internal energy

q = heat gained or lost by the system

w = work done on or by the system

|Target #3: I can describe how the change in internal energy of a system is related to the changes of heat and work between the system and its|

|surroundings. |

A positive “q” means the system has gained heat

A positive “w” means work is done on the system by its surroundings.

When q > 0 and w > 0, then ∆E > 0

When q < 0 and w < 0, then ∆E < 0

When q < 0 and w > 0, then ∆E depends upon the magnitude of q and w

When q > 0 and w < 0, then ∆E depends upon the magnitude of q and w

|Target #4: I can define enthalpy and relate the enthalpy change in a process to the heat added (endothermic) to or lost (exothermic) by the |

|system during the process. |

Enthalpy (H) - heat absorbed or released in a chemical reaction (under constant pressure); the abbreviation qp indicates that pressure is constant; note that ∆H=qp under constant pressure conditions . . . as in coffe cup calorimeters!

Endothermic process: (+∆H) system absorbs heat

2 H2O(g) ( 2 H2(g) + O2(g) ; ∆H = +483.6 kJ (pos. sign)

2 H2O(g) + 483.6 kJ ( 2 H2(g) + O2(g)

Exothermic process: (-∆H) system releases heat

2 H2(g) + O2(g) ( 2 H2O(g) ; ∆H = -483.6 kJ (neg. sign)

2 H2(g) + O2(g) ( 2 H2O(g) + 483.6 kJ

∆Hrxn = ∆H (products) - ∆H (reactants)

|Target #5: I can define the term state function and be provide examples of state functions. |

State function - property of a system which is determined by its state (or conditions, such as temperature & pressure) (see figure 5.10 on page 176); the value of a state function depends only upon the present state of the system, not the path the system took to reach that state

State functions can be recognized by the abbreviation of the property. If the abbreviation is a capitol letter, the property is a state function. If the abbreviation is a lower case letter, then the prperty is not a state function.

Practice: Which of the following are state functions?

a) potential energy e) temperature

b) kinetic energy f) pressure

c) heat g) velocity

d) work h) mass

|Target #6: I can sketch an exothermic and endothermic energy diagram and label the various parts. |

Draw an energy diagram the following reactions. Label the ∆H, energy of reactants and the energy of products.

CaO(s) + CO2(g) ( CaCO3(s) CO2(g) + 2 H2O(l) ( CH4(g) + 2 O2(g)

∆H = -178 kJ ∆H = 890 kJ

|Target #7: I can make stoichiometric calculations based upon a thermochemical equation. |

Consider the following reaction:

2 KClO3(s) ( 2 KCl(aq) + 3 O2(g) ; ∆H = -89.4 kJ

For this reaction, calculate the ∆H for the formation of:

a) 1.24 moles of O2

b) 4.89 grams of KCl

|Target #8: I can differentiate between the terms heat capacity and specific heat. |

| |Definition |Intensive or |Units |

| | |Extensive? | |

|Heat capacity | | | |

|Specific heat (or specific heat | | | |

|capacity) | | | |

|Molar heat capacity | | | |

|Target #9: I can solve problems using q = m•Cp•∆T, calculate any one of the quantities given the other three. |

calorimetry - the measure of heat flow

calorimeter - instrument used to measure heat flow

bomb calorimeter - a special type of calorimeter used for combustion reactions

q = (grams) (specific heat) (∆T)

*When given three of these four terms, you should be able to solve for the fourth!

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example 1: How much heat would it take to raise the temperature of 50.0 kg of water from 25.0oC to 75.0oC?

example 2: Calculate the enthalpy change (in kJ) upon converting 36.0 g of ice at -25.0oC to vapor at 125oC under constant pressure at 1 atm. Use the following constants:

Cp (ice) = 2.09 J/g-K

Cp (H2O-liquid) = 4.18 J/g-K

Cp (water vapor) = 1.84 J/g-K

∆H fusion = 6.01 kJ/mol

∆H vaporization = 40.67 kJ/mol

AP Chemistry Practice Name_________________________

Chapter 5 – Thermochemistry Review Date ____________ Period____

Targets #1-9

Directions: Answer each of the following questions.

1. The KE of a 7.3 kg steel ball traveling at 18.0 m/s is _______ J.

2. One Joule equals . . . .

a) 1 kg m2/s2 b) 2 kg c) 4.184 cal d) 1 g cm/sec e) none of these

3. An 35.0 mg object must be moving at a speed of ___ km/hr in order to posses a KE of 1.20 J.

4. A ______ system can exchange energy with the surroundings but cannot exchange matter

with the surroundings.

5. The change in energy (∆E) of a system that releases 12.4 J of heat and does 4.2 J of work

on the surroundings is ____ J.

6. q has a _____ (+/ -) value when heat is transferred from the surroundings to the system and

work has _____ (+/ -) value when work is done by the system on the surroundings.

7. The sign of ∆E is always negative when q ___ 0 and when w ___ 0.

Fill in the blanks with , or =.

8. Circle all of the following which are state functions?

∆E q heat work temperature ∆PE enthalpy

9. A ____ ∆H (+/ -) corresponds to an endothermic process.

10. The value for ∆H for the following reaction is -72 kJ. How many kJ of heat will be evolved

when 1.0 mol of HBr is formed? H2(g) + Br2(g) ( 2 HBr(g)

11. The value of ∆Ho for the reaction below is -790 kJ. The enthalpy change accompanying the

reaction of 0.95 g of S is _____ kJ. 2 S(s) + 3 O2(g) -----> 2 SO3(g)

12. Draw an energy diagram for the following reaction:

4 Al(s) + 3 O2(g) ( 2 Al2O3(s) + 3351 kJ

13. ∆H = q when ______________________ is held constant. ∆H = ∆E when the

_____________ is held constant.

14. If 73.51 J of heat were added to 25.6 g of water at 25.3oC, what would be the final

temperature of the water?

15. Circle all of the following is/are endothermic processes?

dry ice subliming water evaporating water freezing

16. Under what conditions does ∆H = q?

17. Which of the following is/are true concerning enthalpy?

a) ∆H for a reaction depends upon the states of the reactants and products.

b) ∆H (forward reaction) = -∆H (reverse reaction)

c) The enthalpy of a system is independent of the way the system achieved its current state.

d) Changes in enthalpy must be measured at a constant T, P, and V.

18. 2 Na2O2(s) + 2 H2O(l) ( 4 NaOH(s) + O2(g) ; ∆H = -126 kJ

How much heat is produced when 2.00 mol of NaOH is formed by the above reaction?

Circle true or false for #19-24.

T F 19. Enthalpy is an intensive property.

T F 20. H2O(g) ( H2O(l) is an example of an endothermic process.

T F 21. The enthalpy of a system can be measured.

T F 22. The enthalpy change for a reaction depends upon the state of the reactants and products.

T F 23. ∆Hforward = 1/ ∆ Hreverse

T F 24. The difference between ∆ E and ∆H is the amount of pressure volume work done by the system.

Extra Calorimetry Problems

(Target 9)

|q = m•Cp•∆T |

|Cp of H2O(l) = 4.184 J/g•oC |

|Cp of Hg(l) = 0.14 J/g•oC |

|Cp of Al(s) = 0.90 J/g•oC |

Directions: Solve each of the following. Use the information above to help. Check your answers on the WIKI page!

1. What is the heat capacity (J/oC) of 534 grams of liquid mercury?

2. How many kJ of heat are required to raise the temperature of 450. g of water

from 23.0oC to 85.0oC?

3. Assume you added 0.500 kJ of heat to a 100.0 gram piece of aluminum at 17.0oC.

What is the final temperature of the piece of aluminum?

4. When a 3.88 gram sample of solid NH4NO3 dissolves in 60.0 grams of water in a

coffee cup calorimeter, the temperature drops from 23.0oC to 18.4oC. Calculate

the ∆H (in kJ/mol) for NH4NO3 for the dissolving process. Assume that the

specific heat and density of the solution is the same as pure water.

5. A 75.0 gram piece of aluminum is heated to 95.0oC. The hot piece of aluminum is

placed in a coffee cup calorimeter which holds 100.0 mL of water at 22.0oC.

What will be the final temperature of the water in the calorimeter?

|Target #10: I can describe the difference between a coffe cup calorimeter and a bomb calorimeter. I will also be able to solve simple bomb |

|calorimetry problems. |

| |What is constant? |Important Relationships |

|Coffe cup calorimeters | | |

|Bomb calorimeters | | |

Bomb Calorimetry Problems

The following are examples of the various types of bomb problems that you may see on the test over this unit. These problems are all based on the idea that a calorimeter has a certain heat capacity (usually expressed in J/oC or kJ/oC). The heat capacity of the calorimeter is the amount of heat actually absorbed by the calorimeter itself. The rigid steel container (the “bomb”) usually contains a known mass of the substance to be ignited. The bomb is surrounded by water inside of the calorimeter. The energy change of the reaction inside of the bomb is determined by measuring the increase in the temperature of the water and of the other parts of the calorimeter.

Example #1: A 0.5269-gram sample of octane, C8H18, is placed in a bomb calorimeter known to have a total heat capacity of 11.3 kJ/oC. The octane is ignited and the temperature increase of the calorimeter is 2.25oC. Calculate the amount of energy (in kJ/mol) released by this reaction.

Example #2: Camphor (C10H16O) has a heat of combustion of 5903.6 kJ/mol. When a 0.1204 g sample of camphor is burned in a bomb calorimeter, the temperature increases by 2.28oC. Calculate the heat capacity of the calorimeter.

Example #3: A 1.05 g sample of benzoic acid is combusted completely in a bomb calorimeter. Calculate the change in temperature given that the heat capacity of the calorimeter is 1.80 kJ/K and the heat of combustion of benzoic acid is 26.4 kJ/g.

Example #4: The combustion of methanol (CH3OH) is 239 kJ/mol. A 3.750-g sample of methanol is burned in a bomb calorimeter. The temperature of the calorimeter increased from 20.00oC to 22.90oC. If the calorimeter contained 1.500 kg of water, calculate the heat capacity of the dry calorimeter.

Example #5: 1.048 g of benzene (C6H6) is placed in the bomb compartment of a bomb calorimeter. The water which surrounds the bomb compartment has a mass of 945 g. When the benzene is burned, the temperature increases from 23.640oC to 32.692oC. The heat capacity of the dry calorimeter is 891 J/oC. Find the heat of combustion of benzene (in kJ/g and kJ/mol).

Extra Bomb Calorimetry Problems

1. A quantity of 1.435 g of naphthalene (C10H8) was burned in a constant-volume bomb calorimeter. Consequently, the temperature of the water rose from 20.17oC to 25.84oC. If the quantity of water surrounding the calorimeter was exactly 2000 grams, and the heat capacity of the dry calorimeter was 1.80 kJ/oC, calculate the molar heat of combustion of naphthalene.

2. A 0.828 gram sample of methanol (wood alcohol, CH3OH) is placed in a bomb calorimeter along with enough oxygen gas to ensure complete combustion. The calorimeter contains 1.35 kg of water, and the heat capacity of the calorimeter interior (without the water) is 1.06 kJ/oC. The molar heat of combustion for methanol is 725.9 kJ/mol. Assuming the initial temperature of the water is 23.10oC, what will be the final temperature of the water after combustion?

|Target #11: I can state Hess’s law and use Hess’s Law to solve problems. |

Hess’s Law - if a reaction is carried out in a series of steps, ∆H for the reaction is equal to the sum of the enthalpy changes for the individual steps

Example 1: Find the ∆H for the following reaction using Hess’s Law and the equations that below:

C(graphite) + 2H2(g) ( CH4(g) (overall reaction)

We can calculate the enthalpy change indirectly using Hess’s Law. The following reactions involving C, H2, and CH4 with O2 have been studied and the ∆Horxn values are accurately known:

(a) C(graphite) + O2(g) ------> CO2(g) ∆Horxn = -393.5 kJ

(b) 2 H2(g) + O2(g) --------> 2 H2O(g) ∆Horxn = -571.6 kJ

(c) CH4(g) + 2 O2(g) --------> CO2(g) + 2 H2O(g) ∆Horxn = -890.4 kJ

Since we want to obtain one equation containing only C and H2 as reactants and CH4 as product, we must reverse (c) to get . . .

CO2(g) + 2 H2O(g) -------> CH4(g) + 2 O2(g) ∆Horxn = +890.4 kJ

The next step is to add the equations:

(a) C(graphite) + O2(g) ------> CO2(g) ∆Horxn = -393.5 kJ

(b) 2 H2(g) + O2(g) --------> 2 H2O(g) ∆Horxn = -571.6 kJ

(c) CO2(g) + 2 H2O(g) -------> CH4(g) + 2 O2(g) ∆Horxn = +890.4 kJ

----------------------------------------------------------------------------- ---------------------------

C(graphite) + 2H2(g) --------> CH4(g) ∆Horxn = -74.7 kJ

Example 2: Calculate the ∆H for the following reaction . . .

2 C(graphite) + H2(g) --------> C2H2(g)

Use the following equations . . .

(a) C(graphite) + O2(g) --------> CO2(g) ∆Horxn = -393.5 kJ

(b) H2(g) + 1/2 O2(g) --------> H2O(l) ∆Horxn = -285.8 kJ

(c) 2 C2H2(g) + 5 O2(g) --------> 4 CO2(g) + 2 H2O(l) ∆Horxn = -2598.8 kJ

Solution . . . We want to eliminate O2, CO2, and H2O from the equation. First we reverse (c) in order to get the C2H2(g) on the products side.

(a) C(graphite) + O2(g) --------> CO2(g) ∆Horxn = -393.5 kJ

(b) H2(g) + 1/2 O2(g) --------> H2O(l) ∆Horxn = -285.8 kJ

(c) 4 CO2(g) + 2 H2O(l) -------> 2 C2H2(g) + 5 O2(g) ∆Horxn = +2598.8 kJ

Next, we multiply (a) by 4 and (b) by 2 and carry out the addition:

(a) 4 C(graphite) + 4 O2(g) --------> 4 CO2(g) ∆Horxn = -1574.0 kJ

(b) 2 H2(g) + O2(g) --------> 2 H2O(l) ∆Horxn = -571.6 kJ

(c) 4 CO2(g) + 2 H2O(l) -------> 2 C2H2(g) + 5 O2(g) ∆Horxn = +2598.8 kJ

---------------------------------------------------------------------- ------------------------------

4 C(graphite) + 2 H2(g) --------> 2 C2H2(g) ∆Horxn = 453.2 kJ

or 2 C(graphite) + H2(g) --------> C2H2(g) ∆Horxn = 226.6 kJ

Example 3: Determine the heat of reaction for the synthesis of propane from its elements by using reactions #1-3 below:

3C(s) + 4H2 (g) -------> CH3CH2CH3(g) ∆Horxn = ?????

rxn #1: CH3CH2CH3(g) + 5 O2(g) ----> 3CO2(g) + 4H2O(g) ∆Ho1 = -2,044 kJ

rxn #2: 2H2 (g) + O2(g) ------> 2H2O(g) ∆Ho2 = -483.6 kJ

rxn #3: C(s) + O2(g) -------> CO2(g) ∆Ho3 = -393.5 kJ

|Target #12: I can define the terms standard state and standard heat of formation. I will also be to write a chemical reaction associated |

|with the standard heat of formation. |

Standard enthalpy of formation (heats of formation) = ∆Hof = heat lost or gained when forming one mole of a substance from its elements at standard state conditions (25oC and 1 atm)

Example: Na(s) + ½ Cl2(g) ( NaCl(s) ; ∆Hof = -410.9 kJ/mol

Table 5.3 and Appendix C in your text lists the standard enthalpies of formations of several compounds and elements. A formation reaction is defined as one in which the reactants are elements in their standard states and forming only ONE mole of product. Below are listed 2 examples of formation reactions:

2 Fe(s) + 3/2 O2(g) ( Fe2O3(s) ∆Hof = -822.16 kJ/mol

2 C(graphite) + 3 H2(g) + 1/2 O2(g) ( C2H5OH(l) ∆Hof = -277.7 kJ/mol

Note: The standard enthalpy of formation for the most stable form of any element is zero.

Practice #1: Which of the following reactions is the value for ∆Horxn equal to ∆Hof for the product?

a) H2(g) + ½ O2 (g) ( H2O(l)

b) 2 H2 (g) + O2 (g) ( 2 H2O(l)

c) 2 H2O(l) + O2 (g) ( 2 H2O2 (l)

d) H2O(l) + ½ O2 (g) ( H2O2 (l)

Practice #2: Circle all of the following which have a ∆Hof equal to zero.

Ne(g) Cl2(g) H2O(g) NH3(aq) Fe(s) H2O(l) F2(l)

Practice #3: Write the formation reaction for butanol, C4H9OH.

|Target #13: I can calculate the enthalpy change in a reaction when given the standard enthalpies of formations of each reactant and product.|

Use the following equation when solving enthalpy of formation problems:

∆Horxn = ∑∆Hoproducts - ∑∆Horeactants

Example #1: Consider the following chemical equation and answer the questions that follow:

2 Al(s) + Fe2O3(s) ( 2 Fe(s) + Al2O3(s)

A) Solve for the enthalpy change (∆Ho) for the above reaction using the thermochemical

data from Appendix C.

B) Solve for the enthalpy change (∆Ho) for the reaction given above by using the equations below:

4 Al(s) + 3 O2(g) ( 2 Al2O3(s) ; ∆Ho = -3339.65 kJ

2 Fe(s) + 3/2 O2(g) ( Fe2O3(s) ; ∆Ho = -822.16 kJ

Example #2:

A) Write the formation reactions for C3H8(g), CO2(g), and H2O(l). Look up the ∆Hof for each.

B) Using ∆Horxn = ∑∆Hoproducts - ∑∆Horeactants, calculate the ∆Ho for the burning of 1 mole of

C3H8(g).

C) Now, calculate the ∆Ho for the combustion of propane using Hess’s Law.

2003 AP Thermochemistry Question

In an experiment, a sample of an unknown, pure gaseous hydrocarbon was analyzed. Results showed that the sample contained 6.000 g of carbon and 1.344 g of hydrogen.

(a) Determine the emperical formula of the hydrocarbon.

(b) The density of of the hydrocarbon at 25oC and 1.09 atm is 1.96 g L-1.

(i) Calculate the molar mass of the hydrocarbon.

(ii) Determine the molecular formula of the hydrocarbon.

In another experiement, liquid heptane, C7H16(l), is completely combusted to produce CO2(g) and H2O(l), as represented by the following equation.

C7H16 (l) + 11 O2 (g) ( 7 CO2 (g) + 8 H2O(l)

The heat of combustion, ∆Hocomb, for one mle of C7H16 (l) is -4.85 X 103 kJ.

(c) Using the information in the table below, calculate the ∆Hof for C7H16 (l) in kJ mol-1.

|Compound |∆Hof (kJ mol-1) |

|CO2 (g) |-393.5 |

|H2O(l) |-285.8 |

(d) A 0.0108 mol sample of C7H16 (l) is combusted in a bomb calorimeter.

(i) Calculate the amount of heat released to the calorimeter.

(ii) Given that the total heat capacity of the calorimeter is 9.273 kJ oC-1, calculate the temperature change of the calorimeter.

Practice Thermochemistry Multiple Choice Questions

Directions: Circle the best answer(s) for each question. You may use calculators.

(The specific heat of water is 4.184 J/g(oC)

1. An endothermic reaction causes the surroundings to . . .

a. warm up

b. become acidic

c. condense

d. decrease in temperature

e. release CO2

2. A glass containing 200.0 g of water at 20.oC was placed in a refrigerator. The water loses

11.7 kJ as it cools to a constant temperature. What is the new temperature?

a. 0.013oC b. 4.0oC c. 6.0oC d. 14oC e. 34oC

3. A piece of copper with a mass of 218 grams has a heat capacity of 83.9 J/oC. What is the

specific heat of copper?

a. 0.385 J/g(oC b. 1830 J/g(oC c. 2.60 J/g(oC d. 1.32 J/g(oC e. 0.00454 J/g(oC

4. When 0.7521 g of benzoic acid was burned in a calorimeter containing 1,000 g of water, a

temperature rise of 3.60oC was observed. What is the heat capacity of the bomb

calorimeter, excluding the water? The heat of combustion of benzoic acid is -26.42 kJ/g.

a. 15.87 kJ/oC b. 4.18 kJ/oC c. 5.52 kJ/oC d. 1.34 kJ/oC e. 752.1 kJ/oC

5. Which of the following which is/are ENDOTHERMIC? There may be more than one correct

answer. If so, circle all correct answers.

a. O2(g) + H2 (g) ( 2 H2O(g)

b. H2O(g) ( H2O(l)

c. 3 O2(g) + 2 CH3OH(g) ( 2 CO2 (g) + 2 H2O(g)

d. H2O(s) ( H2O(l)

e. CO2 (s) ( CO2 (g)

6. A 100.0 mL sample of 0.200 M hydrochloric acid is added to 100.0 mL of 0.200 M aqueous

ammonia in a calorimeter whose heat capacity is 480. J/K. The following neutralization

reaction occurs when the two solution are mixed:

HCl(aq) + NH3(aq) ( NH4Cl(aq)

The temperature increase is 2.34oC. Calculate the ∆H per mole of HCl reacted.

a. 154 kJ b. 1.96 kJ c. 485 kJ d. -1.96 kJ e. -154 kJ

7. To which one of the following reactions occuring at 25oC does the symbol ∆Hof of H2SO4(l)

refer?

a. 2 H(g) + S(g) + 4 O(g) ( H2SO4(l)

b. H2(g) + S(g) + 2 O2 (g) ( H2SO4(l)

c. H2SO4(l) ( H2(g) + S(g) + 4 O(g)

d. H2SO4(l) ( 2H(g) + S(s) + 4 O(g)

e. H2 (g) + S(s) + 2 O2 (g) ( H2SO4(l)

8. When 0.560 grams of Na(s) reacts with excess F2(g) to form NaF(s), 13.8 kJ of heat is

evolved at standard state conditions. What is the standard enthalpy of formation, ∆Hof, of

NaF(s)?

a. 24.8 kJ/mol

b. 570. kJ/mol

c. -24.8 kJ/mol

d. -7.80 kJ/mol

e. -570 kJ/mol

9. Ethanol undergoes combustion according to the following equation:

C2H5OH(l) + O2(g) ( CO2(g) + H2O(l) (not balanced)

The standard heat of combustion of ethanol is -1366.8 kJ/mol. Consider the following table

of information:

|Substance |∆Hof (kJ/mol) |

|CO2(g) |-393.5 |

|H2O(l) |-285.8 |

|C2H5OH(l) |? ? ? ? |

What is the standard enthalpy of formation of ethanol, C2H5OH(l)?

a. 3,010 kJ b. -687.6 kJ c. -277.3 kJ d. 687.6 kJ e. 1,367 kJ

10. Calculate the enthalpy change for the reaction . . .

C(graphite) + 2 H2(g) + ½ O2(g) ( CH3OH(l)

Using the following information:

C(graphite) + O2(g) ( CO2(g) ∆H = -393.5 kJ

H2 (g) + ½ O2 (g) ( H2O(l) ∆H = -285.8 kJ

CH3OH(l) + 3/2 O2 (g) ( CO2 (g) + 2 H2O(l) ∆H = -726.4 kJ

a. -1,691.5 kJ b. -238.7 kJ c. 1691.5 kJ d. 47.1 kJ e. -47.1 kJ

11. Consider the following reaction . . .

H2(g) + ½ O2(g) ( H2O(l) ∆Hof = -286 kJ

What is the enthalpy change for the following reaction?

2 H2O(l) ( 2 H2(g) + O2(g) ∆Horeaction = ???

a. ∆Horeaction = -286 kJ

b. ∆Horeaction = +286 kJ

c. ∆Horeaction = -572 kJ

d. ∆Horeaction = +572 kJ

e. ∆Horeaction = -143 kJ

12. The combustion of octane produces heat according to the following equation:

2 C8H18(l) + 25 O2(g) ( 16 CO2(g) + 18 H2O(l) ∆Horxn = -11,020 kJ

What is the heat of combustion per gram of octane?

a. -5,510 MJ/g

b. -96.5 kJ/g

c. -48.2 kJ/g

d. -193 kJ/g

e. -629.2 kJ/g

13. Objects can possess energy as . . .

I. endothermic energy

II. potential energy

III. kinetic energy

a. I only b. II only c. III only d. I and III only e. II and III only

14. At what velocity (m/s) must a 20.0 g object be moving in order to possess a kinetic energy of 1.00 J?

a. 1.00 b. 100 X 102 c. 10.0 d. 1.00 X 103 e. 50.0

15. The internal energy of a system is always increased by . . .

a. adding heat to a system.

b. having the system do work on the surroundings.

c. withdrawing heat from the system.

d. adding heat to the system and having the system do work on the surroundings.

e. a volume compression.

16. Circle the letters of all of the state functions.

a. H b. q c. w d. heat e. temperature

17. Units of heat capacity are ____ and units of specific heat capacity are _____ .

a. J/g(oC and J/K

b. J/mol and J/mol(oC

c. J/mol(oC and J/mol

d. J and J/K

e. J/K and J/g(oC

18. What is the enthalpy change (in kJ) of the chemical reaction that raises the temperature of

250.0 mL of solution having a density of 1.25 g/mL by 7.80oC? (The specific heat of the

solution is 3.74 J/g(K.)

a. -7.43 b. -12.51 c. 8.20 d. -9.12 e. 6.51

19. The British thermal unit (Btu) is commonly used in engineering applications. A Btu is

defined as the amount of heat required to raise the temperature of 1 pound of water by 1oF.

How many joules are there in one Btu? {Hint: 454 g = 1 pound, oC = (5/9)(oF -32o)}

a. 3,415 b. 60.29 c. 1,054 d. 5.120 X 10-3 e. Not enough info given.

20. How should I prepare for the chapter 5 test?

a. Re-read the chapter.

b. Do all of the problem sets and the practice test. Hint: Pick “e” for #20!

c. Study with a partner or group.

d. Try to get plenty of rest and exercise.

e. All of the above.

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