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Golden Valley HS • AP Chemistry
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5 • Thermochemistry Basics
A BLUFFER’S GUIDE
The Law of Conservation of Energy says that
energy cannot be created or destroyed, it is always
conserved. This is also referred to as the First
Law of Thermodynamics.
The system is the part of the universe that is under
study. Everything not part of the system is
considered the surroundings.
An open system can transfer energy and matter to
and from the surroundings. A closed system is
where energy can be transferred to the
surroundings, but matter cannot.
State functions depend only on the difference
between the final and initial state of the system.
The path to the state of the system does not matter.
E = change in energy (state function)
q = heat, enthalpy (not a state function)
w = work (not a state function)
H = change in heat, heat of rxn (state function)
S = change in entropy (state function)
G = change in Gibb’s free energy (state function)
When H or q is positive (+), the reaction is
endothermic, because it is absorbing heat from
the surroundings therefore they feel cool or cold.
When H or q is negative (–), the reaction is
exothermic, releasing heat to the surroundings
and feel warm or hot.
Heat can be defined in calories, which is the
amount of heat needed to raise 1 gram of H2O
1oC. The joule is the SI unit of energy.
1 calorie = 4.184 J
This is also the specific heat of water (4.184
J/goC). Specific heat of other substances is
defined as the energy needed to raise 1 gram of the
substance 1oC. Substances with larger specific
heat values take more energy to change the
temperature. They also hold their heat longer.
Water has a very high specific heat, whereas
metals have low specific heats.
q = mCT
m is the mass, C is the specific heat (sometimes
abbreviated as Cp) and T is the change in
temperature (final temperature – initial
temperature). The temperature can be measured in
Kelvin or Celsius.
Calorimetry is the measurement of heat flow, it is
measured using a calorimeter. These can be fancy
(bomb calorimeter – a sealed metal cup where a
material can be combusted) or simple (coffee cup
calorimeter).
Enthalpies of Reaction or Heat of Reaction
(Hrxn) measures the change in enthalpy (heat) of
the given reaction. It is measured by:

H = Hproducts – Hreactants
This is also known as “final minus initial.” The
heats involved needed to be provided using the
heats of formation in Appendix C. Note: the Hfo
for elements is 0.
Hess’s Law says that if a reaction is carried out in
a series of steps, H for the overall reaction will
equal the sum of the enthalpy changes for the
individual steps. What you do mathematically to
the operation you must also do the Hrxn.
1. If the coefficients of a chemical reaction
are all multiplied by a constant, the Hrxn
is multiplied by that same constant.
2. If two or more reactions are added together
to obtain an overall reaction, the heats of
these reactions are also added to give the
heat of the overall reaction.
3. If a reaction is reversed, then the Hrxn
must also have its sign changed to its
opposite.