16.1Chapter 16 Theories of Energy Changes • MHR 627
Temperature Change and Heat
All physical changes and chemical reactions are accompanied by changes
in energy. These energy changes are crucial to life on Earth. For example,
chemical reactions in your body generate the heat that helps to regulate
your body temperature. Physical changes, such as evaporation, help to
keep your body cool. On a much larger scale, there would be no life on
Earth without the energy from the nuclear reactions that take place in the
Sun.
The study of energy and energy transfer is known as thermodynamics.
Chemists are interested in the branch of thermodynamics known as
thermochemistry: the study of energy involved in chemical reactions.
In order to discuss energy and its interconversions, thermochemists have
agreed on a number of terms and definitions. You will learn about these
terms and definitions over the next few pages. Then you will examine the
energy changes that accompany temperature changes, chemical reactions,
and physical changes.
Studying Energy Changes
The law of conservation of energy states that the total energy of the
universe is constant. In other words, energy can be neither destroyed
nor created. This idea can be expressed by the following equation:
.Euniverse = 0
Energy can, however, be transferred from one substance to another. It can
also be converted into various forms. In order to interpret energy changes,
scientists must clearly define what part of the universe they are dealing
with. The system is defined as the part of the universe that is being studied
and observed. In a chemical reaction, the system is usually made up of the
reactants and products. By contrast, the surroundings are everything else
in the universe. The two equations below show the relationship between
the universe, a system, and the system’s surroundings.
Universe = System + Surroundings
.Euniverse = .Esystem + .Esurroundings = 0
This relationship is known as the first law of thermodynamics.
According to the first law of thermodynamics, any change in the energy
of a system is accompanied by an equal and opposite change in the energy
of the surroundings.
.Esystem = ..Esurroundings
Consider the chemical reaction that is taking place in the flask in
Figure 16.1. A chemist would probably define the system as the contents
of the flask—the reactants and products. Technically, the rest of the
universe is the surroundings. In reality, however, the entire universe
changes very little when the system changes. Therefore, the surroundings
are usually considered to be only the part of the universe that is likely to
be affected by the energy changes of the system. In Figure 16.1, the flask,
the lab bench, the air in the room, and the student all make up the
surroundings. The system is more likely to significantly influence its
immediate surroundings than it is to influence a mountaintop in Japan
(also, technically, part of the surroundings).
In this section, you will
_ identify and describe
the changes to particle
movement that accompany
a change in temperature
_ describe heat as a transfer
of kinetic energy from a
system of higher temperature
to a system of lower
temperature
_ perform calculations
involving heat capacity,
specific heat capacity,
and mass
_ communicate your understanding
of the following
terms: thermodynamics,
thermochemistry, system,
surroundings, first law of
thermodynamics, open
system, closed system,
isolated system, kinetic
energy, potential energy,
joule (J), temperature (T),
heat (q), specific heat capacity
(c), heat capacity (C)
Section Preview/Outcomes
The solution in the
flask is the system. The flask, the
laboratory, and the student are the
surroundings.
Figure 16.1
628 MHR • Unit 7 Thermochemistry
Depending on how they are separated from their surroundings, systems
are defined in three different ways.
. An open system, as its name implies, is open to its surroundings. Both
energy and matter may be exchanged between an open system and its
surroundings. A reaction in an open beaker is an open system.
. In a closed system, matter cannot move between the system and surroundings.
Energy, however, can be transferred between a closed system
and its surroundings. A reaction in a stoppered Erlenmeyer flask is a
closed system.
. An isolated system is completely insulated from the surroundings.
Neither matter nor energy is exchanged between an isolated system and
its surroundings. You will learn more about the importance of isolated
systems in Chapter 17.
Types of Energy
You may recall from earlier science courses that energy is classified into
two fundamental types. These types of energy are:
. kinetic energy — the energy of motion
. potential energy — energy that is stored
If you pick up a rock and lift it several metres above the ground, the rock
gains potential energy. Once you let go of the rock, the rock falls to the
ground as the potential energy is converted to kinetic energy. As you will
learn, energy changes involved in chemical and physical processes fit into
one or both of these two categories.
The SI unit for both kinetic and potential energy is the joule (symbol J).
The joule is derived from other SI units. One joule is equal to 1 kg_m2
s2 _ .
Temperature Change and Heat
Temperature, T, is a measure of the average kinetic energy of the particles
that make up a substance or system. You can think of temperature as a way
of quantifying how hot or cold a substance is, relative to another substance.
Celsius degrees and Kelvin degrees are the same size. The Kelvin scale begins
at absolute zero. This is the temperature at which the particles in a substance have no kinetic
energy. Therefore, Kelvin temperatures are never negative. By contrast, 0°C is set at the
melting point of water. Celsius temperatures can be positive or negative.
Figure 16.2
373.15
K
380
370
360
350
340
330
320
310
300
290
280
270
260
20
10
0
°C 110
100
90
80
70
60
50
40
30
20
10
0
–10
–250
–260
–270
273.15
0.00
100.00
0.00
.273.15