5.1
thermochemistry the study of the
energy changes that accompany
physical or chemical changes in
matter
Changes in Matter and Energy
What happens when matter undergoes change? Clearly, new substances or states are
produced, but energy changes also occur. If chemistry is the study of matter and its
transformations, then thermochemistry is the study of the energy changes that accompany these transformations. Changes that occur in matter may be classified as physical,
chemical, or nuclear, depending on whether a change has occurred in the arrangements
of the molecules, their electronic structure, or the nuclei of the atoms involved (Figure 1).
Whether ice melts, iron rusts, or an isotope used in medical therapy undergoes radioactive decay, changes occur in the energy of chemical substances.
(a)
(b)
Figure 1
Hydrogen may undergo a physical,
chemical, or nuclear change.
(a) Physical: Hydrogen boils at
252°C (or only about 20°C
above absolute zero):
H2( l) → H2(g)
(b) Chemical: Hydrogen is burned
as fuel in the space shuttle’s
main engines:
2 H2(g) O2(g) → 2 H2O( l)
(c) Nuclear: Hydrogen undergoes
nuclear fusion in the Sun, producing helium:
H H → He
thermal energy energy available
from a substance as a result of the
motion of its molecules
chemical system a set of reactants
and products under study, usually
represented by a chemical equation
surroundings all matter around the
system that is capable of absorbing
or releasing thermal energy
298 Chapter 5
(c)
Heat and Energy Changes
Both physical and chemical changes are involved in the operation of an oxyacetylene
torch to weld metals together. A chemical reaction, which involves ethyne (or acetylene)
and oxygen as reactants, produces carbon dioxide gas, water vapour, and considerable
energy. This energy is released to the surroundings as thermal energy, a form of kinetic
energy that results from the motion of molecules. The result is a physical change — the
melting of the metal — when the increased vibration of metal particles causes them to
break out of their ordered solid pattern. When you are studying such transfers of energy,
it is important to distinguish between the substances undergoing a change, called the
chemical system, and the system’s environment, called the surroundings. A system is
often represented by a chemical equation. For the burning of ethyne, the equation is:
2 C2H2(g) 5 O2(g) → 4 CO2(g) 2 H2O(g) energy
The surroundings in this reaction would include anything that could absorb the
thermal energy that has been released, such as metal parts, the air, and the welder’s protective clothing.
When the reaction occurs, heat, q, is transferred between substances. (An object possesses thermal energy but cannot possess heat.) When heat transfers between a system
and its surroundings, measurements of the temperature of the surroundings are used to
classify the change as exothermic or endothermic (Figure 2).
The acetylene torch reaction is clearly an exothermic reaction because heat flows into
the surroundings. Chemical potential energy in the system is converted to heat energy,
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Section 5.1
which is transferred to the surroundings and used to increase the thermal energy of the
molecules of metal and air. Since the molecules in the surroundings have greater kinetic
energy, the temperature of the surroundings increases measurably.
Chemical systems may be further classified. A chemical reaction that produces a gas
in a solution in a beaker is described as an open system, since both energy and matter
can flow into or out of the system. The surroundings include the beaker itself, the surface on which the beaker sits, and the air around the beaker. In the same way, most
explosive reactions are considered to be open systems because it is so difficult to contain
the energy and matter produced. Figure 3 shows an open system. Most calculations of
energy changes involve systems in which careful measurements of mass and temperature
changes are made (Figure 4). These are considered to be isolated systems for the purpose of calculation. However, it is impossible to completely prevent energy from entering
or leaving any system. In reality, the contents of a calorimeter, or of any container that
prevents movement of matter, form a closed system.
Exothermic
Endothermic
system
Figure 2
In exothermic changes, energy is
released from the system, usually
causing an increase in the temperature of the surroundings. In
endothermic changes, energy is
absorbed by the system, usually
causing a decrease in the temperature of the surroundings.
heat amount of energy transferred
between substances
exothermic releasing thermal
energy as heat flows out of the
system
endothermic absorbing thermal
energy as heat flows into the system
Figure 3
A burning marshmallow is an example of an
open system. Gases and energy are free to
flow out of the system.
temperature average kinetic
energy of the particles in a sample
of matter
open system one in which both
matter and energy can move in or
out
isolated system an ideal system in
which neither matter nor energy can
move in or out
closed system one in which
energy can move in or out, but not
matter
Figure 4
A bomb calorimeter is a device in
which a fuel is burned inside an
insulated container to obtain accurate measurements of heat transfer
during chemical reactions. Because
neither mass nor energy can escape,
the chemical system is described as
isolated.
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Thermochemistry 299
Practice
Understanding Concepts
1. Identify each of the following as a physical, chemical, or nuclear change, with reasons
for your choice:
(a) a gas barbecue operating
(b) an ice cube melting in someone’s hand
(c) white gas burning in a camping lantern
(d) wax melting on a hot stove
(e) zinc metal added to an acid solution in a beaker
(f) ice applied to an athletic injury
2. Identify the system and surroundings in each of the examples in the previous question.
3. Identify the following as examples of open or isolated systems and explain your iden-
tification:
(a) gasoline burning in an automobile engine
(b) snow melting on a lawn in the spring
(c) a candle burning on a restaurant table
(d) the addition of baking soda to vinegar in a beaker
(e) a gas barbecue operating
4. A thimbleful of water at 100°C has a higher temperature than a swimming pool full of
water at 20°C, but the pool has more thermal energy than the thimble. Explain.
5. Identify each of the following as an exothermic or endothermic reaction:
(a) hydrogen undergoes nuclear fusion in the Sun to produce helium atoms;
(b) the butane in a lighter burns;
(c) the metal on a safety sprinkler on the ceiling of an office melts when a flame is
brought near it.
Making Connections
6. (a) List five changes that you might encounter outside your school laboratory. Create
a table to classify each change as physical, chemical, or nuclear; endothermic or
exothermic; and occurring in an open or an isolated system.
(b) What are the most commonly encountered types of chemical reactions in terms
of energy flow?
7. The energy content of foods is sometimes stated in “calories” rather than the SI unit of
Figure 5
The fuel in the burner releases heat
energy that is absorbed by the surroundings, which include the
beaker, water, and air.
joules. Physical activity is described as “burning calories”. Research the answers to
the following questions:
(a) What are the relationships among a calorie, a Calorie, and a joule?
(b) Are calories actually burned? Why is this terminology used?
(c) What laboratory methods are used to determine the energy content of foods?
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Measuring Energy Changes: Calorimetry
calorimetry the technological
process of measuring energy
changes in a chemical system
300 Chapter 5
When methane reacts with oxygen in a lab burner, enough heat is transferred to the
surroundings to increase the temperature and even to cause a change of state (Figure 5).
How is this amount of heat measured? The experimental technique is called calorimetry
and it depends on careful measurements of masses and temperature changes. When a fuel
like methane burns, heat is transferred from the chemical system into the surroundings
(which include the water in the beaker). If more heat is transferred, the observed temperature rise in the water is greater. Similarly, given the same amount of heat, a small
amount of water will undergo a greater increase in temperature than a large amount of
water. Finally, different substances vary in their ability to absorb amounts of heat.
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Section 5.1
These three factors — mass (m), temperature change (∆T), and type of substance —
are combined in an equation to represent the quantity of heat (q) transferred:
q mcT
specific heat capacity quantity of
heat required to raise the temperature of a unit mass of a substance
1°C or 1K
Table 1 Specific Heat Capacities
of Substances
where c is the specific heat capacity, the quantity of heat required to raise the temperature of a unit mass (e.g., one gram) of a substance by one degree Celsius or one
kelvin. For example, the specific heat capacity of water is 4.18 J/(g•°C). (Recall that the
SI unit for energy is the joule, J.) Specific heat capacities vary from substance to substance, and even for different states of the same substance (Table 1).
As the equation indicates, the quantity of heat, q, that flows varies directly with the quantity of substance (mass m), the specific heat capacity, c, and the temperature change, ∆T.
Cancelling units in your calculations will help ensure that you have applied the formula correctly. In this book, quantities of heat transferred are calculated as absolute
values by subtracting the lower temperature from the higher temperature.
Calculating Quantity of Heat
Substance
Specific heat
capacity, c
ice
2.01 J/(g•°C)
water
4.18 J/(g•°C)
steam
2.01 J/(g•°C)
aluminum
0.900 J/(g•°C)
iron
0.444 J/(g•°C)
methanol
2.918 J/(g•°C)
SAMPLE problem
When 600 mL of water in an electric kettle is heated from 20°C to 85°C to make a
cup of tea, how much heat flows into the water?
First, use the density formula to calculate the mass of water.
m dV
600 mL
1.00 g/mL
600 g
Use the heat formula, q mcT, to calculate the quantity of heat transferred.
q
?
m 600 g
c
4.18 J/(g•°C) (from Table 1)
T 85°C 20°C 65°C
q mcT
4.18J
600 g 65°C
(g
•° C
)
1.63 105 J or 163 kJ.
163 kJ of heat flows into the water.
Example
What would the final temperature be if 250.0 J of heat were transferred into 10.0 g of
methanol initially at 20.0°C?
Solution
m 10.0 g
c 2.918 J/(g•°C)
T1 20.0 °C
T T2 – T1 T2 – 20.0°C
q 250 J
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Thermochemistry 301
q mcT
q
T mc
250 J
10.0 g
2.918 J/(g
•°C)
8.57°C
T2 20°C 8.57°C
T2 20.0 8.57
T2 28.6°C
The final temperature of the methanol is 28.6°C.
Practice
Answers
9. 506 kJ
10. 38 g
11. 20ºC
12. (a) 15 M
(b) $77
13. (a) 1.0 104 kJ
(b) $55
Understanding Concepts
8. If the same amount of heat were added to individual 1-g samples of water,
methanol, and aluminum, which substance would undergo the greatest temperature change? Explain.
9. There is 1.50 kg of water in a kettle. Calculate the quantity of heat that flows into
the water when it is heated from 18.0°C to 98.7°C.
10. On a mountaineering expedition, a climber heats water from 0°C to 50°C.
Calculate the mass of water that could be warmed by the addition of
8.00 kJ of heat.
11. Aqueous ethylene glycol is commonly used in car radiators as an antifreeze and
coolant. A 50% ethylene glycol solution in a radiator has a specific heat capacity
of 3.5 J/(g•°C). What temperature change would be observed in a solution of
4 kg of ethylene glycol if it absorbs 250 kJ of heat?
12. Solar energy can preheat cold water for domestic hot-water tanks.
(a) What quantity of heat is obtained from solar energy if 100 kg of water is preheated from 10°C to 45°C?
(b) If natural gas costs 0.351¢/MJ, calculate the money saved if the volume of
water in part (a) is heated 1500 times per year.
13. The solar-heated water in the previous question might be heated to the final
temperature in a natural gas water heater.
(a) What quantity of heat flows into 100 L (100 kg) of water heated from 45°C to
70°C?
(b) At 0.351¢/MJ, what is the cost of heating 100 kg of water by this amount,
1500 times per year?
Heat Transfer and Enthalpy Change
Chemical systems have many different forms of energy, both kinetic and potential. These
include the kinetic energies of
• moving electrons within atoms;
• the vibration of atoms connected by chemical bonds; and
• the rotation and translation of molecules that are made up of these atoms.
More importantly, they also include
• the nuclear potential energy of protons and neutrons in atomic nuclei; and
• the electronic potential energy of atoms connected by chemical bonds.
302 Chapter 5
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Section 5.1
Researchers have not yet found a way to measure the sum of all these kinetic and
potential energies of a system. For this reason chemists usually study the enthalpy
change, or the energy absorbed from or released to the surroundings when a system
changes from reactants to products.
An enthalpy change is given the symbol ∆H, pronounced “delta H,” and can be determined from the energy changes of the surroundings. A useful assumption that will be
applied in more detail later in this chapter is that the enthalpy change of the system
equals the quantity of heat that flows from the system to its surroundings, or from the
surroundings to the system (Figure 6). This assumption applies as long as there is no significant production of gas, which is the case in most reactions you will encounter.
This idea is consistent with the law of conservation of energy — energy may be converted from one form to another, or transferred from one set of molecules to another,
but the total energy of the system and its surroundings remains the same.
enthalpy change (H) the difference in enthalpies of reactants and
products during a change
LEARNING
TIP
In searching through references
you may find the terms enthalpy of
reaction, heat of reaction, change
in heat content, enthalpy change,
and H. They all mean the same
thing.
Hsystem qsurroundings
INVESTIGATION 5.1.1
For example, consider the reaction that occurs when zinc metal is added to hydrochloric
acid in a flask:
Zn(s) 2 HCl(aq) → H2(g) ZnCl2(aq)
Some of the chemical potential energy in the system is converted initially to increased
kinetic energy of the products. Eventually, through collisions, this kinetic energy is transferred to particles in the surroundings. The enthalpy change in the system is equal to
the heat released to the surroundings. We can observe this transfer of energy, and can
measure it by recording the increase in temperature of the surroundings (which include
the solvent water molecules, the flask, and the air around the flask). Our calculations of
the heat released will involve the masses of the various substances as well as their temperature change and specific heat capacities.
In order to control variables and allow comparisons, energy changes in chemical systems are measured at standard conditions of temperature and pressure, such as SATP,
before and after the reaction. Under these conditions, the enthalpy change of a chemical
Changes in Kinetic and Potential Energy
high potential energy
Medical Cold Packs (p. 347)
Can you identify the active chemical
in a medical cold pack?
DID YOU
KNOW
?
Setting Hard
The setting of concrete is quite
exothermic, and the rate at which
it sets or cures determines the
hardness of the concrete. If the
concrete sets too quickly (for
example, if the heat of reaction is
not dissipated quickly enough into
the air), the concrete may expand
and crack.
high kinetic energy
Energy
low kinetic energy
low potential energy
Reaction Progress
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Figure 6
In this example of an exothermic
change, the change in potential
energy of the system (H ) equals
the change in kinetic energy of the
surroundings (q). This is consistent
with the law of conservation of
energy.
Thermochemistry 303
physical change a change in the
form of a substance, in which no
chemical bonds are broken
chemical change a change in the
chemical bonds between atoms,
resulting in the rearrangement of
atoms into new substances
nuclear change a change in the
protons or neutrons in an atom,
resulting in the formation of new
atoms
system is the change in the chemical potential energy of the system because the kinetic
energies of the system’s molecules stay constant (for our purposes at this stage).
We can observe enthalpy changes during phase changes, chemical reactions, or nuclear
reactions. Although the magnitudes of the enthalpy changes that accompany these events
vary considerably (Table 2 and Figure 7), the basic concepts of enthalpy change and
heat transfer apply. Notice how much more energy is produced in a nuclear change than
in a chemical change, and in a chemical change than in a physical change.
Table 2 Types of Enthalpy Changes
Physical changes
• Energy is used to overcome or allow intermolecular forces to act.
• Fundamental particles remain unchanged at the molecular level.
1024 J
1021 J
daily solar energy
falling on Earth
1018 J
energy of a strong
earthquake
1015 J
daily electrical
output of hydroelectric
plant
1012 J
1000 tonnes of
coal burned
109 J
106 J
103 J
100 J
1 tonne of TNT exploded
1 kilowatt-hour of
electrical energy
heat released from
combustion of 1 mol
glucose
1 calorie (4.184 J)
• Temperature remains constant during changes of state
(e.g., water vapour sublimes to form frost: H2O(g) → H2O(s) + heat).
• Temperature changes during dissolving of pure solutes
(e.g., potassium chloride dissolves: KCl(s) + heat → KCl(aq)).
• Typical enthalpy changes are in the range H 100 102 kJ/mol.
Chemical changes
• Energy changes overcome the electronic structure and chemical bonds within the particles
(atoms or ions).
• New substances with new chemical bonding are formed
(e.g., combustion of propane in a barbecue: C3H8(g) 5 O2(g) → 3 CO2(g) + 4 H2O(g) heat);
(e.g., calcium reacts with water: Ca(s) 2 H2O(l) → H2(g) Ca(OH)2(aq) heat).
• Typical enthalpy changes are in the range H 102 – 104 kJ/mol.
Nuclear changes
• Energy changes overcome the forces between protons and neutrons in nuclei.
10–3 J
• New atoms, with different numbers of protons or neutrons, are formed
4
234
(e.g., nuclear decay of uranium-238: 238
92 U → 2He 90 Th heat).
10–6 J
• Typical enthalpy changes are in the range H 1010 1012 kJ/mol. The magnitude of the
energy change is a consequence of Einstein’s equation (Figure 8).
10–9 J
heat absorbed during
division of one bacterial cell
10–12 J
energy from fission of
one 235U atom
10–15 J
14. Explain how Hsystem and qsurroundings are different and how they are similar.
Applying Inquiry Skills
average kinetic energy of
a molecule in air at 300 K
Figure 7
Log scale of the enthalpy changes
resulting from a variety of physical,
chemical, and nuclear changes
304
Understanding Concepts
15. How do enthalpy changes of physical, chemical, and nuclear changes compare?
10–18 J
10–21 J
Practice
Chapter 5
16. Design an experiment to determine the identity of an unknown metal, clearly
describing the set of observations that you would make and the calculations that you
would perform (including units), given the following information:
•
The metal is zinc, magnesium, or aluminum, all of which are shiny, silvery metals.
•
These metals react when placed in dilute acid solution.
•
Dilute acid has the same density and specific heat capacity as water.
•
A Chemical Handbook provides values for the heat (in J) released per unit mass
(g) of metal reacting in acid.
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Section 5.1
Figure 8
In Einstein’s famous equation, large
amounts of energy, E, are produced
when a small amount of mass, m, is
destroyed because c, the speed of
light, is such a large value
(3.0 108 m/s).
Section 5.1 Questions
Understanding Concepts
1. For the three states of matter (solid, liquid, and gas), there
are six possible changes of state. Which changes of state
are exothermic? Which are endothermic?
2. What three factors are involved in calculations of the
amount of heat absorbed or released in a chemical reaction?
3. Identify each of the following as a physical, chemical, or
nuclear change, giving reasons for your choice:
(a) gasoline burning in a car engine
(b) water evaporating from a lake
(c) uranium fuel encased in concrete in a reactor
4. Identify the chemical system and the surroundings in each
of the examples in question 3.
5. Identify each of the examples in question 3 as an open or
mole that would be transferred in each of the changes of
state (to the nearest power of ten).
Making Connections
7. The bomb calorimeter is a commonly used laboratory
apparatus. Research and write a brief report describing the
applications of this technology.
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8. Hot packs and cold packs use chemical reactions to
produce or absorb energy. Write a brief report describing
the chemical systems used in these products and their
usefulness.
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an isolated system. Explain your classifications.
6. Describe the chemical system in each of the examples in
question 3. Compare the relative amounts of energy per
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Thermochemistry 305