```Chapter 5 Thermochemistry
Thermodynamics and Thermochemistry
Chapter 5
Thermochemistry:
• Why do chemical reactions occur? One major
driving force for reactions is stability: Substances
that are high in energy are generally less stable and
more reactive, while substances that are lower in
energy are generally more stable and less reactive.
Energy Changes in Reactions
• Thermodynamics is the study of heat and its
transformations.
Chapter Objectives:
• Understand potential and kinetic energy, and the first law of
thermodynamics.
• Thermochemistry is the branch of thermodynamics
that deals with the relationships between chemistry
and energy.
• Understand the concept of enthalpy, and use standard heats of
formation and Hess’s Law to calculate enthalpy changes.
• Learn how to use specific heat to perform calculations involving
energy changes.
C(s) + O2(g)  CO2(g) + heat
Mr. Kevin A. Boudreaux
Angelo State University
CHEM 1411 General Chemistry
Chemistry: The Science in Context (Gilbert, 4th ed, 2015)
www.angelo.edu/faculty/kboudrea
2KClO3(s) + heat  2KCl(s) + 3O2(g)
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Energy, Work, and Heat
• Energy is the capacity to do work or to supply heat.
Energy can be exchanged between objects by some
combination of either heat or work:
Energy = heat + work
DE = q + w
Energy and the
First Law of
Thermodynamics
• Work is done when a force is exerted
through a distance:
Work = force × distance
• Heat is the flow of energy caused by a
temperature difference.
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Kinetic and Potential Energy
• Kinetic energy (EK) is the energy due to the motion
of an object with mass m and velocity v:
EK = ½ mv2
– Thermal energy, the energy associated with the
temperature of an object, is a form of kinetic
energy, because it arises from the vibrations of
the atoms and molecules within the object.
• Potential energy (EP) is energy due to position, or
any other form of “stored” energy. There are several
forms of potential energy:
– Gravitational potential energy
– Mechanical potential energy
– Chemical potential energy (stored in chemical
bonds)
Figure 5.1 Heat is transferred from the hotter block to
the colder block until thermal equilibrium is reached.
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Kinetic and Potential Energy
• Potential energy increases when things that attract
each other are separated or when things that repel
each other are moved closer.
• Potential energy decreases when things that attract
each other are moved closer, or when things that
repel each other are separated.
• According to the law of conservation of energy,
energy cannot be created or destroyed, but can be
converted from one form to another.
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Chapter 5 Thermochemistry
Kinetic and Potential Energy
Units of Energy
• Water falling in a waterfall exchanges gravitational
potential energy for kinetic energy as it falls faster
and faster, but the energy is never destroyed.
• From the expression for kinetic energy it can be seen
that energy has units of kg·m2/s2 (kg m2 s-2)
• The SI unit of energy is the joule (J):
1 J = 1 kg m2 s-2
[James Prescott
Joule, 1818-1889]
– Since joules are fairly small units, kilojoules (kJ)
• A calorie (cal) is the amount of energy needed to
raise the temperature of 1 g of water by 1°C.
1 cal = 4.184 J (exact)
high EP
low EK
decreasing EP
increasing EK
low EP
high EK
EK converted to
thermal energy
and sound
• The nutritional unit Calorie (Cal) is actually a
kilocalorie (kcal):
1 Cal = 1 kcal = 1000 cal = 4184 J
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Units of Energy
Examples: Kinetic Energy
• The watt (W) is a unit of power, the rate at which
energy is used:
1. What is the kinetic energy (in kJ) of a 2300. lb car
moving at 55.0 mi/h?
1 W = 1 J s-1
• Kilowatt-hours (kWh) are the unit that is usually
reported on electrical bills:
1 kWh = 3.60×106 J
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Internal Energy, E, and State Functions
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State Functions
• The internal energy (E) of a system is the sum of
all the kinetic and potential energies for every
particle in the system.
Change in altitude is a state function: it only depends on the difference
between the initial and final values, not on the path that the climbers take.
• The internal energy is a state function, which is
dependent only on the present state of the system,
and not on the pathway by which it got to that state.
– Some examples of state functions include energy
(and many other thermodynamic terms), pressure,
volume, altitude, distance, etc.
• An energy change in a system can occur by many
pathways (different combinations of heat and
work), but no matter what the combination, DE is
always the same — the amount of the energy change
does not depend on how the change takes place.
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Figure 5.2
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Chapter 5 Thermochemistry
State Function — A Chemical Example
C8 H18 (l) 
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State Functions and Reversibility
• Another characteristic of state functions is their
reversibility.
O2 (g)  8CO2 (g)  9H 2O(g)
initial state ( Einitial )
final state ( Efinal )
– Climbing back down Mount Everest, your final
position is now identical to your initial position,
and the value of the change in altitude is 0 miles.
– The overall change in a state function is zero
when the system returns to its original condition.
• For a nonstate function, the overall change is not
zero even if the path returns the system to its
original condition.
– For instance, the money spent on a trip or time
expended during traveling does not reappear.
• For a given change, DE (q + w) is constant, even
though the specific values of q and w can vary.
• Heat and work are NOT state functions, but DE is.
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The System and the Surroundings
The System and the Surroundings
• A system is the particular set of substances (a
chemical reaction, a phase change, etc.) that we
focus on in an experiment.
• An isolated system does not exchange energy or
matter with its surroundings (e.g., an ideal Thermos
bottle).
• The surroundings include everything else outside
the system — the flask, the room, the building, etc.
• A closed system exchanges energy but not matter
with its surroundings (e.g., a cup of soup with a lid)
• An open system exchanges energy and matter with
its surroundings (e.g., a bowl of soup that is open to
the air, which can evaporate, or have things added).
surroundings
surroundings
surroundings
Solution of A + B
boundary
system
energy flow
across boundary
surroundings
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Fig.
5.11
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Exothermic and Endothermic Processes
The First Law: Conservation of Energy
• If the system loses energy, the surroundings gain
energy, and vice versa. Energy changes are usually
considered from the point of view of the system.
• The law of conservation of energy, also known as
the first law of thermodynamics, states that energy
cannot be created or destroyed; it can only be
converted from one form into another.
– An exothermic process is one where energy
flows out of the system.
• In other words, the total energy of the universe is a
constant.
– An endothermic process is one where energy
flows into the system.
DEuniverse = DEsystem + DEsurroundings = 0
• A process that is exothermic in one direction (e.g.,
from gas to liquid) is endothermic in the other
direction (from liquid to gas)
Figure 5.12
• Potential and kinetic energy are interconvertible, but
for a particular system the total amount of potential
and kinetic energy is a constant.
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Chapter 5 Thermochemistry
Change in Internal Energy
Energy Flow: Sign Conventions
• For an isolated system, with no energy flowing in or
out of the system, the internal energy is a constant.
• By convention, energy changes are measured from
the point of view of the system.
– First Law of Thermodynamics (restated): The
total internal energy of an isolated system is
constant.
– Any energy that flows from the system to the
surroundings has a negative sign: Einitial
Efinal < Einitial
• It is impossible to completely isolate a reaction from
its surroundings, but it is possible to measure the
change in the internal energy of the system, DE, as
energy flows into the system from the surroundings
or flows from the system into the surroundings.
DE < 0
energy
Efinal
– Any energy that flows to the system from the
surroundings has a positive sign:
Efinal > Einitial
DE = Efinal - Einitial
DE > 0
DE = Eproducts - Ereactants
Efinal
energy
Einitial
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Energy Flow: Sign Conventions
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Examples: Energy Flow
• Energy exchanges between the system and
surroundings take place through a combination of
heat (q) and work (w):
2. When gasoline burns in a car engine, the heat
released causes the products CO2 and H2O to
expand, which pushes the pistons outward. Excess
heat is removed by the car’s cooling system. If the
expanding gases do 451 J of work on the pistons and
the system loses 325 J to the surroundings as heat,
calculate the change in energy in J, kJ, and kcal.
DE = q + w
• Energy going into the system, whether it’s heat
energy or work energy, has a positive sign, and
energy leaving the system, whether it’s heat energy
or work energy, has a negative sign:
q is (+) Heat is absorbed by the system
q is (-)
Heat is released by the system
w is (+) Work is done on the system
w is (-)
Work is done by the system
Answer: -776 J, -0.776 kJ, -0.185 kcal
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Work
Pressure-Volume Work
• Work (w) is defined as the force (F) that produces
the movement of an object through a distance (d):
Work = force × distance
w=Fd
• When propane, C3H8, is burned, 7 moles of gas are
produced for every 6 moles of reactants:
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
– If the reaction takes place in a container with a
movable piston, the greater volume of gas product
will force the piston outward against the pressure
of the atmosphere, thus doing work on the piston.
• Work also has units of J, kJ, cal, kcal, Cal, etc.
• The two most important types of chemical work are:
– the electrical work done by moving charged
particles (which is important in electrochemistry).
– the expansion work done as a result of
a volume change in a system, particularly
from an expanding or contracting gas.
This is also known as pressure-volume
work, or PV work.
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– Pressure-volume work is equal to the pressure
(P) times the change in volume (DV); the sign is
negative because work is being done by the
system (work energy is leaving the system):
w = - PDV
• The units of PV work are L·atm; 1 L·atm = 101.3 J.
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Chapter 5 Thermochemistry
Expansion Work
Expansion Work
w  F d ; P  F/A
For a cylinder, V  Ah
w  PDV
w  -PDV
w = - PDV
• If the gas expands, DV is positive, and the work term
will have a negative sign (work energy is leaving the
system).
• If the gas contracts, DV is negative, and the work term
will have a positive sign (work energy is entering the
system).
• If there is no change in volume, DV = 0, and there is
no work done. (This occurs in reactions in which
there is no change in the number of moles of gas.)
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Examples: PV Work
3. Calculate the work done (in kJ) during a reaction in
which the volume expands from 12.0 L to 14.5 L
against an external pressure of 5.00 atm.
Energy Changes at Constant Pressure; Enthalpy
• If a process is carried out at constant volume (DV =
0), the energy change of the reaction appears entirely
as heat and not work (E = q - PDV).
– Most reactions are carried out in open vessels at
constant pressure, with the volume capable of
changing freely.
Enthalpy
– In these cases, DV  0, and the energy change
may be due to both heat transfer and PV work.
• In order to eliminate the contribution from PV work,
a quantity called enthalpy, H, is defined as internal
energy (E) plus the product of pressure and volume:
H = E + PV
DH = DE + PDV
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Chapter 5 Thermochemistry
Energy Changes at Constant Pressure; Enthalpy
• While DE is a measure of the entire energy change
of a system (heat and work), DH is a measure of
only the heat exchanged at constant pressure.
• By doing a little algebra, we can see the relationship
between DH and energy:
DErxn = q + w
Enthalpy vs. Energy
DErxn = q - PDV
• Most reactions do take place at constant (i.e.,
atmospheric) pressure, so DH is a good measure of
energy changes in most thermodynamic systems.
DH = DE + PDV
DH = (q - PDV) + PDV
– In reactions that do not involve gases, volume
changes are approximately zero, and DH  DE.
DH = qP
• Therefore, the change in enthalpy equals the heat
evolved at constant pressure, qP.
– When the number of moles of gas doesn’t change,
DV = 0, and DH = DE.
• The enthalpy change, DH, is also called the heat of
reaction for the process. It is also a state function:
– Even in reactions in which the number of moles
of gas does change, qP is usually much larger than
PDV, so DH  DE.
DH = Hfinal - Hinitial = Hproducts - Hreactants
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Examples: Enthalpy vs. Energy
Exothermic vs. Endothermic Processes
• A chemical reaction can either release or absorb
heat energy:
1. The reaction of nitrogen with hydrogen to make
ammonia has DH° = -92.2 kJ
N2(g) + 3H2(g)  2NH3(g); DH° = -92.2 kJ
– An endothermic process absorbs heat energy
and results in an increase in the enthalpy of the
system (the products have more enthalpy than the
reactants).
What is the value of DE (in kJ) if the reaction is
carried out at a constant pressure of 40.0 atm and the
volume change is -1.12 L?
– An exothermic process releases heat energy and
results in a decrease in the enthalpy of the system
(the products have less enthalpy than the
reactants).
Products
Hfinal < Hinitial
DH < 0
exothermic
Hfinal > Hinitial
DH > 0
endothermic
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Enthalpy Diagrams
Enthalpy/
Energy
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Where Does the Energy Come From?
• When an exothermic reaction occurs, the chemical
potential energy that used to be stored in the
chemical bonds in the reactants.
Reactants
– The higher energy bonds break and lower energy
bonds form, and the system goes from a higher
potential energy to a lower potential energy.
Reactants
Products
Heat is absorbed
endothermic
DH > 0
Heat is released
exothermic
DH < 0
– The energy difference is usually emitted in the
form of heat or light.
• When an endothermic reaction occurs, the energy
that is absorbed by the system allows new bonds
with higher energies to be formed, and the potential
energy of the system is raised.
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Chapter 5 Thermochemistry
Thermochemical Equations
Examples: Thermochemical Equations
• A thermochemical equation is a balanced equation
that also states the heat of reaction (DH).
2. How much heat (in kJ) is evolved when 5.00 g of
aluminum reacts with a stoichiometric amount of
Fe2O3 in the thermite reaction? (WE 8.3)
• The reaction of 1 mole of N2 and 3 moles of H2 to
form 2 moles of NH3 releases 92.38 kJ of energy:
2Al(s) + Fe2O3(s)  2Fe(s) + Al2O3(s); DH° = -852 kJ
N2(g) + 3H2(g)  2NH3(g); DH = -92.38 kJ
– This DH value is true only when the coefficients
are taken to mean the number of actual moles
they express for that particular reaction, with the
reagents in those particular physical states.
– For other amounts of reactants or products, the
energy change of the reaction can be calculated
by using the DH term in the same way that the
coefficients are used.
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MOV: Thermite
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Examples: Thermochemical Equations
3. If aluminum is produced by the thermal
decomposition of bauxite (mostly Al2O3), how many
grams of Al can form when 1.000  103 kJ of heat is
transferred?
Al2O3(s)  2Al(s) + 3/2O2(g); DH = +1676 kJ
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Heat and Temperature
• Heat and temperature are not the same thing!
– Heat is the exchange of thermal energy between a
system and its surroundings. Temperature is a
measure of the thermal energy of a sample.
Heat Capacity
and
Heating Curves
– A fresh cup of coffee is hotter than a swimming
pool, but a swimming pool holds more heat than
the coffee. If you pour a 200°F cup of coffee into
a 100°F pool, the pool won’t become 150°F.
• Heat flows from higher temperatures to lower
temperatures, until the two regions reach the same
temperature, and are at thermal equilibrium.
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• All forms of energy can be converted quantitatively
into heat, so it is possible to use heat as a measure of
energy for any chemical or physical process.
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Chapter 5 Thermochemistry
Heat Capacity
Specific Heat
• The thermal properties of a substance are those
that describe its ability to absorb or release heat
without changing chemically.
• The specific heat (c, or specific heat capacity, Cs in
book) of an object, is the quantity of heat required to
change the temperature of 1 gram of a substance by
1 °C (or K):
• The heat capacity (C) of an object is the amount of
heat (q) required to change its temperature by a
given amount (DT):
C 
q
DT
Cs 
• The molar heat capacity (Cm), is the quantity of
heat required to change the temperature of 1 mole of
a substance by 1 °C (or K):
– Heat capacity has units of J/°C (or J/K), and is an
extensive property, depending on the sample size.
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Specific Heats of Some Common Substances
Specific Heat
Capacity, Cs
(J g-1 °C-1)
Substance
0.128
Ethanol
2.42
Gold
0.128
Water (l)
4.184
Platinum
0.133
Water (s)
2.03
Tin
0.228
Water (g)
1.865
Silver
0.235
Carbon (s)
0.71
Copper
0.385
Glass (Pyrex)
0.75
Zinc
0.388
Granite
0.79
Nickel
0.444
Sand
0.84
Iron
0.449
Titanium
0.522
Aluminum
0.903
Magnesium
1.02
Cm 
q
moles DT
q  Cm moles DT
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Heat of Fusion & Heat of Vaporization
• During a phase change, there is no change in
temperature as heat is added, because both physical
states are present during a phase change (there is an
equilibrium between the two phases).
Specific Heat
Capacity, Cs
(J g-1 °C-1)
• The enthalpy change when 1 mole of a solid melts
(fuses) into a liquid is called the heat of fusion,
DHfus. The enthalpy change during fusion is given by
q = n DHfus
• The enthalpy change when 1 mole of a liquid
vaporizes into a gas is called the heat of
vaporization, DHvap. The enthalpy change during
vaporization is given by
q = n DHvap
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Examples: Specific Heat
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Examples: Specific Heat
1. If a gold ring with a mass of 10.00 g changes in
temperature from 25.00°C to 28.00°C, how much
energy in joules has it absorbed? How much energy
is absorbed by 10.00 g of H2O under the same
conditions?
Answer: qAu = 3.87 J, qH2O = 125 J
q  Cs m DT
– Specific heat has units of J / g °C, and is an
intensive property, which is independent of the
sample size.
q  C DT
DT  Tf - Ti
Substance
q
m DT
2. What is the specific heat of the (fictional) element
dilithium if it takes 2.51 kJ to raise the temperature
of 42.0 g of dilithium by 10.0°C?
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Chapter 5 Thermochemistry
Examples: Specific Heat
Examples: Specific Heat
3. What is the final temperature of a 100.0 g sample of
water initially at 25.0°C after 25.1 kJ of heat energy
is added? (The specific heat of water is 4.18 J/g°C.)
4. How much energy (in J) is absorbed when 10.0 g of
ice melts to form liquid water? (The heat of fusion
for water is 6.01 kJ/mol.)
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Heating-Cooling Curves
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Heating Curve for Water
• A heating-cooling curve shows the changes that
occur when heat is added or removed at constant rate
from a sample.
• During a phase change, a change in heat occurs at a
constant temperature (q = n DH°phase change).
• Within a phase, a change in heat is accompanied by
a change in temperature (q = Cs m DT).
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Figure 5.25
Heating Curve for Water
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Examples: Heating Curves
• Segment 1: Solid ice warms from -25°C to 0°C.
5. How much energy (in kJ) is needed to heat 270. g of
ice from -18.0°C to liquid water at 85.0°C? (DHvap
= 40.67 kJ/mol, DHfus = 6.01 kJ/mol, the specific
heat of ice is 2.03 J g-1 °C-1, the specific heat of
liquid water is 4.184 J g-1 °C-1, and the specific heat
of steam is 1.865 J g-1 °C-1.)
q = Cs, ice m DT
• Segment 2: Solid ice melts into liquid water at 0°C.
q = n DH°fus
• Segment 3: Liquid water warms from 0°C to 100°C.
q = Cs, liquid m DT
• Segment 4: Liquid water boils into steam at 100°C.
q = n DH°vap
• Segment 5: Steam warms from 100°C to 125°C.
q = Cs, steam m DT
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Chapter 5 Thermochemistry
Constant-Volume Calorimetry: Measuring DE
• For any chemical reaction, the energy change is the
sum of the heat and work energies:
DErxn = q + w = q - PDV
Calorimetry:
Measuring Heat and
Work
• Under conditions of constant volume, w=0, and the
energy change of the reaction is dependent only on
heat evolved at constant volume, qV:
DErxn = qV
• The heat change — and therefore the energy change
— in a chemical reaction occurring at constant
volume can be measured in a bomb calorimeter, in
which a sample is sealed in a “bomb” surrounded by
a water bath, in which the temperature change can
be measured precisely.
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A Bomb Calorimeter
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Examples: Constant-Volume Calorimetry
• The amount of heat
absorbed by the water in
the calorimeter is equal to
the energy that is released
by the reaction (but
opposite in sign).
1. When 1.010 g of sucrose (C12H22O11) undergoes
combustion in a bomb calorimeter, the temperature
rises from 24.92°C to 28.33°C. Find DErxn for the
combustion of sucrose in kJ/mol sucrose. The heat
capacity of the bomb calorimeter (determined in a
separate experiment) is 4.90 kJ/°C. (You can ignore
the heat capacity of the small sample of sucrose
because it is negligible compared to the heat
capacity of the calorimeter)
qcal  Ccal DT
qrxn  qcal
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Figure 5.30
Constant-Pressure Calorimetry: Measuring DH
Examples: Constant-Pressure Calorimetry
• Energy changes at constant pressure are measured
in an insulated vessel with a stirrer, thermometer,
and a loose-fitting lid that keeps the contents at
atmospheric pressure. The data obtained allows us
to calculate DH. (In undergraduate labs, this device
is approximated by a coffee cup.)
2. A 25.64 g sample of a solid was heated in a test tube
to 100.00°C in boiling water and carefully added to
a coffee-cup calorimeter containing 50.00 g water.
The water temperature increased from 25.10°C to
28.49°C. What is the specific heat capacity of the
solid? (Assume all the heat is gained by the water.)
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qsoln  Cs, soln msoln DT
qrxn  qsoln
Figure 5.29
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Chapter 5 Thermochemistry
Examples: Constant-Pressure Calorimetry
3. The reaction of hydrochloric acid and sodium
hydroxide is exothermic. The equation is
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
In one experiment, a student placed 50.0 mL of 1.00
M HCl at 25.50°C in a coffee cup calorimeter. To
this was added 50.0 mL of 1.00 M NaOH solution
also at 25.50°C. The mixture was stirred, and the
temperature quickly increased to a maximum of
32.40°C. What is the energy evolved in kilojoules
per mole of HCl? (Because the solutions are
relatively dilute, we can assume that their specific
heats are close to that of water, 4.18 J g-1 °C-1, and
that their densities are 1.00 g mL-1.)
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Manipulating Thermochemical Equations
• If a chemical equation is multiplied by some factor,
then DHrxn is multiplied by the same factor.
2H2(g) + O2(g)  2H2O(g); DH = -483.6 kJ
4H2(g) + 2O2(g)  4H2O(g); DH = -967.2 kJ
Hess’s Law and
Heats of Formation
H2(g) + ½O2(g)  H2O(g); DH = -241.8 kJ
• If a chemical equation is reversed, the sign of DHrxn
is reversed. (This is equivalent to multiplying the
equation by -1.)
2H2O(g)  2H2(g) + O2(g); DH = +483.6 kJ
H2O(g)  H2(g) + ½O2(g); DH = +241.8 kJ
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Examples: Thermochemical Equations
Two Paths to CO2
• We can imagine two paths leading from 1 mole of
carbon and oxygen (O2) to 1 mole of carbon dioxide:
1. Given the following thermochemical equation
N2(g) + 3H2(g)  2NH3(g); DHrxn = -92.38 kJ
What is the DHrxn for the reactions shown below?
One-Step Path:
a. 2NH3(g)  N2(g) + 3H2(g); DH = +92.38 kJ
C(s) + O2(g)  CO2(g); DH° = -393.5 kJ
b. 2N2(g) + 6H2(g)  4NH3(g); DH = -184.8 kJ
Two-Step Path:
c. 1/2N2(g) + 3/2H2(g)  NH3(g); DH = -46.19 kJ
C(s) + 1/2O2(g)  CO(g); DH° = -110.5 kJ
d. 4NH3(g)  2N2(g) + 6H2(g); DH = +184.8 kJ
CO(g) + 1/2O2(g)  CO2(g); DH° = -283.0 kJ
e. 6NH3(g)  3N2(g) + 9H2(g); DH = +277.1 kJ
f. 2NH3(aq)  N2(g) + 3H2(g); DH = no relation!
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Chapter 5 Thermochemistry
Hess’s Law
Enthalpy Diagram for the Formation of CO2
• Since DH is a state function, the path which leads
from the starting materials to the products should be
irrelevant to the overall energy change.
C(s) + O2(g)
DH° = -110.5 kJ
CO(g) + 1/2O2(g)
DH° = -393.5 kJ
A  B  C; DH  - 100 kJ
C  D; DH   30 kJ
A  B  D; DH  - 70 kJ
DH° = -283.0 kJ
CO2(g)
DH
• This is generalized into Hess’s Law: if a reaction
can be expressed as the sum of a series of steps, then
DH for the overall reaction is the sum of the DH’s of
the individual steps.
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Examples: Hess’s Law
• We can use this property to determine the DH’s of
reactions which are difficult to measure directly, by
writing the reaction as the sum of reactions with
known DH’s.
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Examples: Hess’s Law
2. Calculate the value of DH° for the hypothetical
reaction
3. Hydrogen peroxide, H2O2, decomposes into water
and oxygen by the following equation.
2A  3D
H2O2(l)  H2O(l) + 1/2O2(g)
Use the following equations to determine the value
of DH°.
Use the following equations to determine the value
of DH° for the decomposition of hydrogen peroxide.
A  B
DH° = +10 kJ
H2(g) + O2(g)  H2O2(l); DH° = –188 kJ
3C  2B
DH° = -40 kJ
H2(g) + 1/2O2(g)  H2O(l); DH° = –286 kJ
D  C
DH° = -20 kJ
69
Examples: Hess’s Law
70
Examples: Hess’s Law
5. Use Hess’s law to calculate the value of DH° for the
reaction of tungsten with carbon:
4. Carbon monoxide is often used in metallurgy to
remove oxygen from metal oxides and thereby give
the free metal. The thermochemical equation for the
reaction of CO with iron(III) oxide, Fe2O3, is
W(s) + C(s)  WC(s)
Use the following thermochemical equations:
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g); DH°=-26.7 kJ
2W(s) + 3O2(g)  2WO3(s); DH° = -1680.6 kJ
Use this equation and the equation for the
combustion of carbon monoxide,
C(s) + O2(g)  CO2(g); DH° = -393.5 kJ
CO(g) + 1/2O2(g)  CO2(g); DH° = –283.0 kJ
2WC(s) + 5O2(g)  2WO3(s) + 2CO2(g); DH° = -2391.6 kJ
to calculate the value of ∆H° for the following
reaction.
2Fe(s) + 3/2O2(g)  Fe2O3(s)
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72
Chapter 5 Thermochemistry
Thermodynamic Standard States
Some Important Types of Enthalpy Change
• The enthalpy change is the amount of heat released
or absorbed when reactants are converted to
products at the same temperature and in the molar
amounts represented by coefficients in the balanced
equation. The physical states must also be specified,
since a substance in the gas phase has a different
heat content from one in the liquid phase.
• Heat of fusion (DHfus) — enthalpy change when 1
mole of a substance melts.
H2O(s)  H2O(l); DH = 6.01 kJ/mol at 0°C
• Heat of vaporization (DHvap) — enthalpy change
when 1 mole of a substance vaporizes.
H2O(l)  H2O(g); DH = 40.7 kJ/mol at 100°C
• In order to compare thermodynamic measurements
from different reactions, a set of conditions called
the thermodynamic standard state is defined: the
most stable form of a substance at 1 atm pressure
and 25°C, 1 M for all substances in solution.
• Heat of sublimation (DHsubl) — enthalpy change
when 1 mole of a substance sublimes (converts
directly from the solid phase to the vapor phase).
• Heat of combustion (DHc) — enthalpy change
when 1 mole of a substance combines with O2 in a
combustion reaction.
• When DH is measured with all substances in their
standard states, it is called the standard enthalpy
(heat) of reaction, DH°.
C4H10(l) +
73
Formula
• The standard enthalpy (heat) of formation, DHf°,
is the enthalpy change for the formation of 1 mole of
a substance in its standard state from its constituent
elements in their standard states.
Determining DH° from DHf°
DHf° (kJ/mol)
Formula
-277.6
Nitrogen
74
DHf° (kJ/mol)
111.9
C3H8(g)
-103.85
N2(g)
0
Br2(l)
0
C3H6O(l, acetone)
-248.4
NH3(g)
-45.9
HBr(g)
-36.3
C3H8O(l, isopropanol)
-318.1
NH4NO3(s)
-365.6
C6H6(l)
49.1
NO(g)
91.3
81.6
Ca(s)
0
C6H12O6(s, glucose)
-1273.3
N2O(g)
CaO(s)
-634.9
C12H22O11(s, sucrose)
-2226.1
Oxygen
CaCO3(s)
-1207.6
Chlorine
Carbon
75
Formula
C2H5OH(l)
Br(g)
Calcium
• The elements must be in their standard states at
25°C. That is, elements which are diatomic must be
in the diatomic form; physical states must be
reasonable for that element at 25°C, etc.
O2(g)
0
Cl(g)
121.3
O3(g)
142.7
C(s, graphite)
0
Cl2(g)
0
H2O(g)
-241.8
C(s, diamond)
1.88
HCl(g)
-92.3
H2O(l)
-285.8
CO(g)
-110.5
Fluorine
CO2(g)
-393.5
F(g)
79.38
Na(s)
0
CH4(g)
-74.6
F2(g)
0
Na(g)
107.5
CH3OH(l)
-238.6
HF(g)
-273.3
NaCl(s)
-411.2
C2H2(g)
227.4
Hydrogen
Na2CO3(s)
-1130.7
C2H4(g)
52.4
H(g)
218.0
NaHCO3(s)
-950.8
C2H6(g)
-84.68
H2(g)
0
Sodium
Appendix A4.3on pages APP-15 through APP-21 lists thermodynamic properties of more substances.
76
Examples: Standard Heats of Formation
6. Do the following reactions represent a DHf°
equation?
• The standard enthalpy of reaction can be calculated
by subtracting the sum of the heats of formation of
all reactants from the sum of the heats of formation
appropriate coefficient in the balanced equation:
CO(g) + 1/2O2(g)  CO2(g); DH° = –283.0 kJ
nr reactants  np products
 np DH f (products) -
DHf° (kJ/mol)
Bromine
• The most stable form of any element in its standard
state has DHf° = 0 kJ. This establishes a baseline
from which to calculate the DHf° for other
compounds.
 
DH rxn
 4CO2(g) + 5H2O(l)
Standard Enthalpies of Formation
Standard Enthalpy of Formation (DHf°)
• In most cases, DHf° for a compound is negative (i.e.,
a compound is usually more stable than its
elements.)
13/ O (g)
2 2
2H(g) + O(g)  H2O(l); DH° = -971.1 kJ
 nr DH f (reactants)
2H2(g) + O2(g)  2H2O(l); DH° = –571.8 kJ
• This process is really an application of Hess’s Law.
H2(g) + 1/2O2(g)  H2O(l); DH° = –285.9 kJ
• Thus, to calculate the DH° for any reaction, we write
the balanced equation, and look up the DHf°’s of all
of the reactants and products.
77
78
Chapter 5 Thermochemistry
Examples: Standard Heats of Formation
Examples: Standard Heats of Formation
7. What equation must be used to represent the
formation of NaHCO3(s) when we want to include
its value of DHf°?
8. Some chefs keep baking soda, NaHCO3, handy to
put out grease fires. When thrown on the fire, baking
soda partly smothers the fire and the heat
decomposes it to give CO2, which further smothers
the flame. The equation for the decomposition of
NaHCO3 is
2NaHCO3(s)  Na2CO3(s) + H2O(l) + CO2(g)
Use the DHf° data in Appendix A4.3 to calculate the
DH° for this reaction in kilojoules.
79