Chapter 6 Thermochemistry: Energy Flow and Chemical Change
Except for nuclear reactions where mass and energy significantly interchange,
chemical reactions obey the laws of conservation of matter and conservation of energy
separately within the limits of measurement. What this means for beginning chemistry
students is that atoms are conserved in chemical reactions and energy changes can be
summed up in an energy cycle.
Major Concepts to Know:
 Students need to be able to distinguish the various forms that energy takes in an
“ensemble” of atoms or molecules. Thermal energy is in the random movement of
atoms and molecules (translation, vibration, and rotation) and chemical energy is
stored within the structure of substances in chemical bonds. Chemical reactions
rearrange atoms in molecules and therefore change the chemical energy stored in
bonds. Energy is absorbed, released, or converted to other forms of energy in the
process.

Most chemical reactions involve the absorbing or releasing of heat. Many students
have misconceptions about heat. Temperature is not heat, but a measure of the
thermal energy of a substance. Heat is the transfer of thermal energy between two
bodies or objects. Generally, heat flow is described as moving from the hot object to
the cooler object. An easy way to see if students misunderstand the concept is to ask
students to touch two objects at the same temperature—like glass and steel. They will
say they have a different temperature even though they can look at a thermometer and
see the temperatures are the same! Why? Because the objects each have a different
heat capacity and ability to transfer heat to your hand.

Thermochemistry is the name given to the study of heat exchange in chemical
reactions. Before reactions can be studied, it is vital that basic vocabulary of
reactions and systems is reviewed. Students appreciate that chemical reaction
equations MUST be balanced for any thermal considerations to work. Students must
understand the difference between open, closed, and isolated systems. It is helpful
(though strictly not correct) if students see heat as a reactant or product in
endothermic and exothermic reactions for later on when studying equilibrium shifts:
Exothermic reactions give off heat, so heat is a product; endothermic reactions absorb
heat, so heat is a reactant.

On the AP test, students will see this information presented in a variety of ways and
they need to practice to avoid confusion. They may be told simply that the reaction is
exothermic or endothermic or be given the information in the form of ΔH values.
Strictly, H is the change only in enthalpy, but since other energy changes are usually
small, this suffices for most considerations. A convention has been established that
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energy lost by a system has a negative sign, so for an exothermic reaction, ΔH has a
“–”value; for endothermic the ΔH value would be “+”.

Students should recognize that temperature, volume, pressure, and energy are all state
functions, and the only important values are the initial and final values, regardless of
the steps in between.

The only law of thermodynamics in this chapter is the first law—energy can be
converted from one form to another but not destroyed. Energy change can be
measured using the equation ΔE = Eproducts – Ereactants. If the reaction is exothermic, the
surroundings will heat up from the released energy. If the reaction is endothermic,
the surroundings will cool down from transferring energy into the reacting system.
The total energy change is the sum of the enthalpy change and any work done. So
E = H – PV at constant pressure.

E = H – PV can also be related to work where work can be defined as force
multiplied by distance (w = Fd). In chemistry, for example, this is most easily seen in
a reaction with gases. Work is done by the expansion of a gas in a cylinder, so the
equation w = – PΔV is more useful. Work also has sign conventions students should
know: “–” when work is done by the system on the surroundings (the gas is
expanding and pushing a syringe plunger outward) and “+” when work is done by the
surroundings on the system (a syringe plunger moves inward as a cooling gas shrinks
in volume). Since PV is often in L.atm units, students may need to convert to the
energy unit of joules by using 1L.atm = 101.3 J. Note that if the reaction has no net
change in moles of gas as it proceeds (Δn = 0), no volume change will occur and no
work will be done.

Work and heat are not state functions; they represent transferred quantities only. The
state function enthalpy is the thermodynamic quantity used to describe the heat
content at constant volume and has the symbol H. ΔH then can be used to describe the
change in heat content in chemical reactions shown by the equation: ΔH = Hproducts –
Hreactants.

Students should recognize when solids melt to liquids and when liquids vaporize into
gases or solids sublime. ΔH is “+” as heat is being absorbed, and when each reverses
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(gases undergo deposition to a solid or condensation to a liquid and a liquid freezes
into a solid), the ΔH is “–”.

Students should also know to look at thermochemical equations and recognize what
the ΔH means and how to manipulate it. For example, if an equation is multiplied by
2, the ΔH is also doubled; if an equation is reversed, the sign of ΔH is reversed.

Three terms that students should understand are calorimetry, specific heat (s), and
heat capacity (C). Students often get confused between specific heat and heat
capacity. Specific heat is the amount of energy to raise 1 gram 1°Celsius, whereas
heat capacity is the energy required to raise a given sample 1°Celsius; therefore,
specific heat is intensive (not dependent on the amount) whereas heat capacity is
extensive (dependent on the amount).

A calorimeter is a closed system to measure heat changes. In many problems, by
observing what happens to water in the system, calculations can be done with the
water to indicate what happened to the system. Using the equation q = msΔt, where
m is the mass, s is the specific heat (for water 4.184 J/g. °C), and t is the temperature,
the energy of the system q can be calculated. This can then be related to how much a
mole would release.
Coffee Cup Calorimeter

If a reaction is done in a bomb calorimeter (a constant volume system), the total heat
released or absorbed is the sum of the water and the container (bomb calorimeter).
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To calculate the heat for the bomb calorimeter, the heat capacity must be known. The
equation is qsystem = qcal + qrxn, where qcal = CcalΔt and qrxn(H2O) = mΔts. Remind
students to add the two numbers—they must be in the same units.
A Comb Calorimeter

Another way to obtain the enthalpy of a reaction is to use thermodynamic data. It is
vital for students to grasp the definitions of “formation” and “standard states.” Using
standard ΔH°f at 1 atm and 25°C (in the back of the text), we can calculate the ΔH of
a reaction using ΔH°rxn = Σ ΔH°f(products) – Σ ΔH°f(reactants). Students should recognize
for problems the ΔH°f of elements (in a standard state) is 0, by definition. Many
times, students will think a problem can’t be solved because an element’s value isn’t
given.

An indirect method to determine ΔH can also be used with Hess’s law, and the same
principles of combining reactions used throughout chemistry. If two equations can be
added together to make the desired equation, the ΔH values of the two equations can
also be added. If an equation must be reversed to make the desired equation, reverse
the sign of ΔH, and if an equation must be doubled, then double the ΔH value, etc.

Encourage students to write out all the steps in a Hess’s law problem so they can be
sure the cycle of reactions is complete in order to find the missing value. For
example, see discussion in Chapter 4 about how an ionic solid dissolves by hydration.
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
Students should also be reminded that tables in the back of the book are for standard
states, per mole of substance, and at 298K. In other conditions, especially at other
temperatures and amounts, the prediction may be different.
Vocabulary to Know:
 Calorimetry
 Endothermic process
 Enthalpy (H)
 Exothermic process
 First law of thermodynamics
 Heat
 Heat capacity (C)
 Law of conservation of energy
 Specific heat (c)
 Standard enthalpy of formation
 Standard heat of fusion
 Standard heat of vaporization
 Standard enthalpy of reaction
 Standard state
 Thermal energy
 Thermochemical equation
 Thermodynamics
 Work
Math Skills Students Must Know:
 ΔE = Eproducts – Ereactants
 E = H – PV
 w = – PΔV
 q = msΔt = CΔt
 ΔH°rxn = Σ ΔH°f(products) – Σ ΔH°f(reactants)
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Suggested Problems:
 Forms of Energy and Their Interconversion: 8, 9, 97
 Enthalpy: Heats of Reaction and Chemical Change: 16, 18, 19, 21
 Calorimetry: Laboratory Measurements of Heats of Reactions: 29, 30, 33, 34, 35, 36,
37, 39, 40, 41, 42, 85
 Stoichiometry of Thermochemical Equations: 50, 51, 52, 53, 54
 Hess’s Law of Heat Summation: 63, 64, 101, 102
 Standard Heats of Reaction: 75, 76, 77, 78
Suggested Demonstrations or Labs:
 Cooper, Melanie M. “Project 12: Hot and Cold,” Cooperative Chemistry Lab Manual
(McGraw-Hill, 2006).
 Paradis, Jeffrey A. “Introducation to Thermochemistry: Using a Calorimeter,” Hands
on Chemistry Laboratory Manual (McGraw-Hill, 2006).
 Paradis, Jeffrey A. “Calorimetry: Nutrition in a Nutshell,” Hands on Chemistry
Laboratory Manual (McGraw-Hill, 2006).
 Paradis, Jeffrey A. “Hess’s Law: A Study of the Combustion of Magnesium,” Hands
on Chemistry Laboratory Manual (McGraw-Hill, 2006).
 Shakhashiri, Bassam Z. “1.3 Endothermic Reactions of Hydrated Barium Hydroxide
and Ammonium Salts,” Chemical Demonstrations: A Handbook for Teachers of
Chemistry, Volume 2 (The University of Wisconsin Press, 1989).
Questions
1. What is energy?
2. What are some of the types of energy?
a.
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b.
c.
d.
3. What is true when energy is transferred from one object to another?
4. Explain each of the following terms.
a. System
b. Surroundings
c. Internal energy
5. What is the relationship between the change in energy of a system and the energy
change in the surroundings?
6. What is heat?
7. Which way is heat moving in each of the following types of reactions and what is the
sign of each?
a. Exothermic
b. Endothermic
8. Explain the conditions when work is positive and when work is negative.
a. How can gases perform work?
9. What is the first law of thermodynamics?
a.
How can change in energy be measured?
b. What is a joule?
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c. What is a calorie?
d. What are state functions?
e. Is the change in energy a state function? Explain your answer.
f. Is work or heat a state function? Explain your answer.
10. What is enthalpy?
a. What is the symbol for enthalpy?
b. In what three types of chemical reactions is work not considered when calculating
energy change?
c. What is the equation to solve for enthalpy?
d. Why is a balanced equation important when solving for enthalpy?
e. What is an exothermic reaction?
f. What is an endothermic reaction?
g. Explain the difference between the heat of combustion, heat of formation, heat of
fusion, and the heat of vaporization.
11. What is the difference between specific heat, heat capacity, and molar heat capacity?
a.
What is specific heat and what is its symbol?
b. Give the equation for specific heat.
c. What is the unit for specific heat?
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d. What is heat capacity and what is its symbol?
e. What is the equation for heat capacity?
f. What is the unit for heat capacity?
12. What is a calorimeter?
a. What is assumed in solving equations with a calorimeter?
b. What is a bomb calorimeter?
c. Why must the heat capacity of the bomb be accounted for?
d. What is the difference between a constant volume and a constant pressure
calorimeter?
13. What is the difference between a thermochemical equation and a chemical equation?
14. What is the relationship between sign and magnitude of the ∆H between the forward
reaction and the backward reverse reaction?
15. Explain the concept of Hess’s law of heat summation.
a. If you double an equation, what do you do to the heat?
b. If you reverse an equation, what do you do to the sign?
16. What is standard enthalpy of formation and what is the symbol?
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a. What is standard state for heats of formation?
b. What is the enthalpy of any element in its standard state?
c. What is the equation for solving for enthalpy using heats of formation?
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