The Foundations of Chemistry

the atomic or molecular level, we can think of each of these as either kinetic or potential
energy.
The chemical energy in a fuel or food comes from potential energy stored in atoms
due to their arrangements in the molecules. This stored chemical energy can be released
when compounds undergo chemical changes, such as those that occur in combustion and
metabolism. Reactions that release energy in the form of heat are called exothermicreac-
tions.
Combustion reactions of fossil fuels are familiar examples of exothermic reactions.
Hydrocarbons — including methane, the main component of natural gas, and octane, a
minor component of gasoline — undergo combustion with an excess of O 2 to yield CO 2
and H 2 O. These reactions release heat energy. The amounts of heat energy released at
constant pressure are shown for the reactions of one mole of methane and of two moles
of octane.

CH 4 (g)2O 2 (g)88nCO 2 (g)2H 2 O(
)890 kJ

2C 8 H 18 (
)25O 2 (g)88n16CO 2 (g)18H 2 O(
)1.090 104 kJ

In such reactions, the total energy of the products is lower than that of the reactants by
the amount of energy released, most of which is heat. Some initial activation (e.g., by heat)
is needed to get these reactions started. This is shown for CH 4 in Figure 15-1. This acti-
vation energy plus890 kJ is released as one mole of CO 2 (g) and two moles of H 2 O(
) are
formed. A process that absorbs energy from its surroundings is called endothermic.One
such process is shown in Figure 15-2.
Energy changes accompany physical changes, too (Chapter 13). For example, the
melting of one mole of ice at 0°C at constant pressure must be accompanied by the absorp-
tion of 6.02 kJ of energy.

H 2 O(s)6.02 kJ88nH 2 O(
)

This tells us that the total energy of the water is raised by 6.02 kJ in the form of heat
during the phase change.

A hydrocarbon is a binary compound

of only hydrogen and carbon.

Hydrocarbons may be gaseous, liquid,

or solid. All burn.

The amount of heat shown in such an

equation always refers to the reaction

for the number of moles of reactants

and products specified by the

coefficients. We call this one mole of

reaction.It is important to specify the

physical states of all substances,

because different physical states have

different energy contents.

15-1 The First Law of Thermodynamics 593

See the Saunders Interactive

General Chemistry CD-ROM,

Screen 6.12, Energy Changes in

Chemical Processes.

Figure 15-1 The difference

between the potential energy of the

reactants — one mole of CH 4 (g) and

two moles of O 2 (g) — and that of

the products — one mole of CO 2 (g)

and two moles of H 2 O(
) — is the

amount of heat evolved in this

exothermicreaction at constant

pressure. For this reaction, it is

890 kJ/mol of reaction. In this

chapter, we see how to measure the

heat absorbed or released and how

to calculate it from other known

heat changes. Some initial activation,

for example by heat, is needed to get

the reaction started. In the absence

of such activation energy, a mixture

of CH 4 and O 2 can be kept at room

temperature for a long time without

reacting. For an endothermicreaction,

the final level is higher than the

initial level.

CH 4 (g)  2O 2 (g)

Potential Energy

CO 2 (g)  2H 2 O(
)

Reactants Products

Progress of reaction

Heat evolved  890 kJ

Amount of energy

needed to activate

the reaction