General Chemistry:
An Integrated Approach
Hill, Petrucci, 4th Edition
Chapter 6
Thermochemistry
Mark P. Heitz
State University of New York at Brockport
© 2005, Prentice Hall, Inc.
Energy
Literally means “work within,” however no object
contains work
Energy refers to the capacity to do work
– that is, to move or displace matter
2 basic types of energy:
– Potential (possibility of doing
work because of composition or
position)
– Kinetic (moving objects doing
work)
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Energy
Potential Energy – in a
gravitational field
(= position)
Kinetic Energy –
energy of motion
PE = mgh
m = mass (kg)
g = gravity
constant (m s–2)
h = height (m)
KE = 1/2mv2
m = mass (kg)
v = velocity (m s–1)
v2 = (m2 s–2)
units are kg m2 s–2  J
units are kg m2 s–2  J
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Work
Work is the product of the force in the direction of
motion and the distance the object is moved
Work = force × distance  energy (J)
Collisions in the real world are not perfectly elastic
Energy transfer occurs as a result
of inelastic collisions
e.g., the ball loses height
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Thermochemistry
System
Surroundings
Universe
Surroundings
Thermochemistry is the study of energy changes that
occur during chemical reactions
Surroundings
Focus is on heat
and matter
transfer between
the system ...
and the
surroundings
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Thermochemistry
Types of systems one can study:
Matter
Matter
Energy
OPEN
Energy
Matter Matter
Energy
Matter Matter
Energy
CLOSED
Energy Energy
ISOLATED
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Internal Energy
Internal Energy (U) is the total energy contained
within the system, partly as kinetic energy and
partly as potential energy
Kinetic involves three types of molecular motion ...
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Internal Energy
Internal Energy (U) is the total energy contained
within the system, partly as kinetic energy and
partly as potential energy
Potential energy involves
intramolecular interactions ...
and intermolecular
interactions ...
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Heat (q)
Heat is energy transfer resulting from thermal
differences between the system and surroundings
“flows” spontaneously
from higher T 
lower T
“flow” ceases at
thermal equilibrium
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Heat Transfer Mechanism
Illustrated
Inelastic
molecular
collisions
are
responsible
for heat
transfer
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Heat Transfer Illustrated
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Work (w)
Work is an energy transfer between a system
and its surroundings
Recall from gas laws … the product PV = energy
Pressure–volume
work is the work
of compression
(or expansion) of
a gas
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Calculating Work (w)
PV work is calculated as follows:
w = –PDV
Sign conventions: think
SYSTEM WORK
from the perspective of the
system
If work is done by the
system, the system loses
energy equal to –w
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Calculating Work (w)
SYSTEM WORK
Expansion is an example of
work done by the system—the
weight above the gas is lifted
compression (or
expansion) of a
gas
ExpansionWork
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Calculating Work (w)
SYSTEM
WORK
If work is done on the system,
the system gains energy equal
to +w
compression (or
expansion) of a
gas
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States of a System
The state of a system refers to its exact condition,
determined by the kinds and amounts of
matter present, the structure of this matter at
the molecular level, and the prevailing pressure
and temperature
Example: internal energy (U) is a
function of the state of the system ...
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State Functions
A state function is a
property that has a unique
value that depends only on
the present state of a
system and not on how the
state was reached, nor on
the history of the system
DU = Uf – Ui
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First Law of Thermodynamics
The Law of Conservation of Energy states that in a
physical or chemical change, energy can be exchanged
between a system and its surroundings, but no energy can
be created or destroyed
The change in U is related to the energy
exchanges that occur as heat (q) and work
(w)
The First Law: DU = q + w
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First Law – Sign Conventions
Energy entering a system carries a positive sign: if
heat is absorbed by the system, q > 0. If work is done
on a system, w > 0
Energy leaving a system carries a negative sign: if
heat is given off by the system, q < 0. If work is done
by a system, w < 0
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Heats of Reaction
The heat of reaction (qrxn) is the quantity of heat
exchanged between the system and its surroundings
Examples – for exothermic reactions,
in isolated systems, system T 
in non-isolated systems, heat is given off to
the surroundings, i.e., q < 0
– endothermic reactions,
in isolated systems, system T 
in non-isolated systems, heat is absorbed
from the surroundings, i.e., q > 0
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Conceptualizing an Exothermic Reaction
Surroundings are at 25 °C
25 °C
Typical situation:
some heat is released
to the surroundings,
some heat is absorbed
by the solution.
Hypothetical situation: all heat
is instantly released to the
surroundings. Heat = qrxn
32.2 °C
35.4 °C
In an isolated system, all heat is
absorbed by the solution.
Maximum temperature rise.
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Internal Energy Changes
w = –PDV and DU = q + w
For systems where the reaction is carried out at
constant volume, DV = 0 and DU = qV
All the thermal energy produced
by conversion from chemical
energy is released as heat
Because the reaction is
exothermic, both qV and
DU are negative
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Internal Energy Changes
w = –PDV and DU = q + w
For systems where the reaction is carried out at
constant pressure, DU = qP – PDV or qP = DU + PDV
Most of the thermal energy is
released as heat, but some is work
used to expand the system against
the surroundings
The quantity of heat liberated is
somewhat less than in the constantvolume case
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Example 6.2
The internal energy of a fixed quantity of an ideal gas depends only on its temperature. If
a sample of an ideal gas is allowed to expand against a constant pressure at a constant
temperature, (a) what is ∆U for the gas? (b) Does the gas do work? (c) Is any heat
exchanged with the surroundings?
Analysis and Conclusions
(a) Because the expansion occurs at a constant temperature, the expanded gas (state 2)
is at a lower pressure than the compressed gas (state 1) but the temperature is
unchanged. Because the internal energy of the ideal gas depends only on the
temperature, U2 = U1 and ∆U = U2 – U1 = 0.
(b) The gas does work in expanding against the confining pressure, P. The pressure–
volume work is w = –P∆V, as was illustrated in Figure 6.8. The work is negative
because it is done by the system.
(c) The work done by the gas represents energy leaving the system. If this were the only
energy exchange between the system and its surroundings, the internal energy of the
system would decrease, and so would the temperature. However, because the
temperature remains constant, the internal energy does not change. This means that
the gas must absorb enough heat from the surroundings to compensate for the work
that it does in expanding: q = –w. And, according to the first law of thermodynamics,
∆U = q + w = –w + w = 0.
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Example 6.2 continued
Exercise 6.2A
In an adiabatic process, a system is thermally insulated from its surroundings so that
there is no exchange of heat (q = 0). If an ideal gas undergoes an adiabatic expansion
against a constant pressure, (a) does the gas do work? (b) Does the internal energy of the
gas increase, decrease, or remain unchanged? (c) What happens to the temperature?
a) yes w = - PΔV
b) ΔU = q + w = 0 + w = w = - PΔV since ΔV >0, ΔU< 0
c) For an ideal gas U ~ T so if U decreases, T decreases
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Enthalpy
Most heats of reaction are measured at constant
pressure … it is useful to have a function equal to qP
Enthalpy (H) is the sum of the internal energy and the
pressure–volume product of a system
qP = DH = DU + PDV
Enthalpy is an extensive property (depends on how
much of the substance is present)
Enthalpy is a state function. U, P, and V are all state
functions, therefore H must be a state function also
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Enthalpy Diagrams
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Reversing a Reaction
• DH changes sign when a process is reversed.
• Therefore, a cyclic process has the value DH = 0.
Same magnitude; different signs.
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Using DH
Values are measured experimentally
Negative values indicate exothermic reactions
Positive values indicate endothermic reactions
Changes sign when a process is reversed.
Therefore, a cyclic process has the value DH = 0
For problem-solving, one can view heat being
absorbed in an endothermic reaction as being like
a reactant and heat being evolved in an
exothermic reaction as being like a product
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Calorimetry
• We measure heat flow using calorimetry.
• A calorimeter is a device used to make this
measurement.
• A “coffee cup” calorimeter may be used for
measuring heat involving solutions.
A “bomb” calorimeter is used to
find heat of combustion; the
“bomb” contains oxygen and a
sample of the material to be burned.
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Calorimetry Relationships
The heat capacity (C) of a system is the quantity of
heat required to change the temperature of the
system by 1 oC
calculated from C = q/DT
units of J oC–1 or J K–1
Specific heat is the heat capacity of a one-gram sample
Specific heat = C/m = q/mDT
units of J g–1 oC–1 or J g–1 K–1
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Specific Heats
Molar heat capacity is the product of specific heat
times the molar mass of a substance
units are J mol–1 K–1
A useful form of the specific heat equation is:
q = m CDT
If DT > 0, then q > 0 and heat is gained by the system
If DT < 0, then q < 0 and heat is lost by the system
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Specific Heats
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Hess’s Law of
Constant Heat Summation
The heat of a reaction is constant, regardless of the
number of steps in the process
DHoverall = S DH’s of individual reactions
When it is necessary to reverse a chemical
equation, change the sign of DH for that reaction
When multiplying equation coefficients, multiply
values of DH for that reaction
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An Enthalpy Diagram
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Standard State Conditions
The standard state of a solid or liquid substance is
the pure element or compound at 1 atm pressure and
the temperature of interest
Gaseous standard state is the “ideal gas” at 1 atm
pressure and the temperature of interest
e.g., at 1 atm, 25 oC standard state for Hg is liquid, C
is solid, water is liquid, He is gas
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Standard Enthalpies
The standard enthalpy of reaction (DHo) is the
enthalpy change for a reaction in which the reactants
in their standard states yield products in their
standard states
The standard enthalpy of formation (DHof) of a
substance is the enthalpy change that occurs in the
formation of 1 mol of the substance from its
elements when both products and reactants are in
their standard states
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Standard Enthalpies
o
of Formation at 25 C
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Calculations Based on
Standard Enthalpies of Formation
General Expression:
DHo = Snp × DHof (products) – Snr × DHof (reactants)
Each coefficient is multiplied by the standard enthalpy
of formation for that substance
The sum of numbers for the reactants is subtracted
from the sum of numbers for the products
With organic compounds, the measured DHof is
often the standard enthalpy of combustion DHocomb
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Standard Enthalpies of Formation
of Ions in Aqueous Solution
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Combustion Fuels
Fossil Fuels: Coal, Natural Gas, and Petroleum
A fuel is a substance that
burns with the release of heat
These fossil fuels were formed
over a period of millions of
years from organic matter that
became buried and compressed
under mud and water
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Respiration Foods
Foods: Fuels for the Body
The three principal classes of foods are fats,
proteins, and carbohydrates
1 Food Calorie (Cal) is equal to 1000 cal
(or 1 kcal)
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Summary of Concepts
• Thermochemistry concerns energy changes in
physical processes or chemical reactions
• Thermochemical ideas include the notion of a
system and its surroundings; the concepts of
kinetic energy, potential energy, and internal
energy; and the distinction between two types of
energy exchanges: heat (q) and work (w)
• Internal energy (U) is a function of state
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Summary of Concepts
• Enthalpy (H) is a function based on internal
energy, but modified for use with constantpressure processes
• The first law of thermodynamics relates the heat
and work exchanged between a system and its
surroundings to changes in the internal energy of a
system
• A calorimeter is used to measure quantities of heat
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Summary of Concepts
• The concepts of standard state, a standard enthalpy
change, and a standard enthalpy of formation are
important in thermochemical calculations
• Some practical applications of thermochemistry
deal with the heats of combustion of fossil fuels
and the energy content of foods
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