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Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Hmwk pg. 188-196, # 7, 12, 15, 19, 21, 23, 29, 33,
41, 45, 51, 53, 59, 61, 67, 71, 77, 83, 97, 103
Chapter 5
Thermochemistry
John D. Bookstaver
St. Charles Community College
St. Peters, MO
© 2006, Prentice Hall, Inc.
Thermochemistry
Energy
•  _______________________________.
Ø _____: Energy used to cause an object
that has mass to move.
Ø _____: Energy used to cause the
temperature of an object to rise.
Thermochemistry
Potential Energy
Energy an object possesses by virtue of its
position or chemical composition.
Thermochemistry
Kinetic Energy
Energy an object possesses by virtue of its
motion.
Thermochemistry
Units of Energy
•  The SI unit of energy is the joule (J).
kg m2
1 J = 1 ⎯⎯
s2
•  An older, non-SI unit is still in widespread use: The
calorie (cal).
1 cal = 4.184 J
•  Calories that we eat are actually __________ are
abbreviated as _____
Thermochemistry
System and Surroundings
•  The _____ includes
the molecules we want
to study (here, the
hydrogen and oxygen
molecules).
•  The _____ are
everything else (here,
the cylinder and
piston).
Thermochemistry
Work
•  Energy used to
move an object over
some distance.
•  w = F ž d,
Thermochemistry
Heat
•  Energy can also be
________________.
•  Heat flows from
_____ objects to
_____ objects.
•  Recall, cold is
simply a
________________.
Thermochemistry
Transferal of Energy
a)  The potential energy of this ball of
clay is ________ when it is moved
from the ground to the top of the wall.
b)  As the ball falls, its potential energy is
converted to _____ energy.
c)  When it hits the ground, its kinetic
energy falls to _____ (since it is no
longer moving); some of the energy
does work on the ball, the rest is
________________.
Thermochemistry
First Law of Thermodynamics
•  Energy ________________________________.
•  In other words, the total energy of the universe is
a constant; if the system loses energy, it must be
gained by the surroundings, and vice versa.
Use Fig. 5.5
Thermochemistry
Internal Energy
The internal energy of a system is the sum of all
kinetic and potential energies of all components
of the system; we call it E.
Use Fig. 5.5
Thermochemistry
Internal Energy
By definition, the change in internal energy, ΔE,
is the final energy of the system minus the initial
energy of the system:
Use Fig. 5.5
Thermochemistry
Changes in Internal Energy
•  If ΔE > 0, Efinal > Einitial
Ø Therefore, the system
absorbed energy from
the surroundings.
Ø This energy change is
called ________.
Thermochemistry
Changes in Internal Energy
•  If ΔE < 0, Efinal < Einitial
Ø Therefore, the system
released energy to the
surroundings.
Ø This energy change is
called ________.
Thermochemistry
Changes in Internal Energy
•  When energy is
exchanged between
the system and the
surroundings, it is
exchanged as either
heat (q) or work (w).
•  That is
Thermochemistry
ΔE, q, w, and Their Signs
Thermochemistry
Exchange of Heat between
System and Surroundings
•  When heat is
absorbed by the
system from the
surroundings, the
process is
________.
Thermochemistry
Exchange of Heat between
System and Surroundings
When heat is released
by the system to the
surroundings, the
process is
________.
Thermochemistry
State Functions
Usually we have no way of knowing the
internal energy of a system; finding that value
is simply too complex a problem.
Thermochemistry
State Functions
•  However, we do know that the internal energy
of a system is independent of the path by
which the system achieved that state.
Ø In the system below, the water could have reached
room temperature from either direction.
Thermochemistry
State Functions
•  Therefore, internal energy is a ____________.
•  It depends only on ______________________,
not on the path by which the system arrived at
that state.
•  And so, ΔE depends only on Einitial and Efinal.
Thermochemistry
State Functions
•  However, q and w are
_____ state functions.
•  Whether the battery is
shorted out or is
discharged by running
the fan, its ΔE is the
same.
Ø But q and w are different
in the two cases.
Thermochemistry
State Functions
Yes!
Your bank account
balance
No!
The distance traveled
from Penncrest to
K o P mall.
Thermochemistry
Work
When a process
occurs in an
__________,
commonly the only
work done is a change
in volume of a gas
pushing on the
surroundings (or being
pushed on by the
surroundings). Thermochemistry
Work
We can measure the work done by the gas if
the reaction is done in a vessel that has been
fitted with a piston.
Thermochemistry
Enthalpy
•  If a process takes place at ____________
(as the majority of processes we study do)
and the only work done is this
______________, we can account for
heat flow during the process by
measuring the ________ of the system.
•  Enthalpy is the internal energy plus the
product of pressure and volume:
Thermochemistry
Enthalpy
•  When the system changes at constant
pressure, the change in enthalpy, ΔH, is
ΔH = Δ(E + PV)
•  This can be written
ΔH = ΔE + PΔV
Thermochemistry
Enthalpy
•  Since ΔE = q + w and w = −PΔV, we
can substitute these into the enthalpy
expression:
ΔH = ΔE + PΔV
•  So, at constant pressure the change in
enthalpy is the heat gained or lost.
Thermochemistry
Endothermicity and
Exothermicity
•  A process is
endothermic when
ΔH is _____.
•  A process is
exothermic when
ΔH is _____.
Thermochemistry
Endo or Exo? State whether ΔH
is positive or negative
• 
• 
• 
• 
• 
The combustion of fuel in a car engine
Salting winter roads
Making snow on the ski trails
Skiing down the trails
Snow-covered trails melt in the spring
time.
•  Placing an ice pack on an injury
Thermochemistry
Enthalpies of Reaction
The ________ in
enthalpy, ΔH, is the
enthalpy of the
products minus the
enthalpy of the
reactants:
Thermochemistry
Enthalpies of Reaction
This quantity, ΔH, is called the enthalpy of
reaction, or the _______________.
Thermochemistry
The Truth about Enthalpy
1.  Enthalpy is an ________ property.
2.  ΔH for a reaction in the forward
direction is equal in size, but
__________, to ΔH for the reverse
reaction.
3.  ΔH for a reaction depends on the state
of the products and the state of the
reactants.
Thermochemistry
Enthalpy Calculation
Sucrose is oxdized to carbon dioxide and
water. The enthalpy change can be
measured in the laboratory as shown
C12H22O11 + 12O2 → 12CO2 + 11H2O ΔH = -5645 kJ
What is the enthalpy change for the
oxidation of 5.00 g of sugar?
Thermochemistry
Calorimetry
Since we cannot
know the
________________
________________,
we measure ΔH
through calorimetry,
the measurement of
________________.
Thermochemistry
Heat Capacity and Specific Heat
•  The amount of energy required to raise
the temperature of a substance by 1 K
(1°C) is its _________________.
•  We define _________________ as the
amount of energy required to raise the
temperature of 1 g of a substance by 1
K.
Thermochemistry
Heat Capacity and Specific Heat
Specific heat, then, is
heat transferred
Specific heat =
mass × temperature change
Thermochemistry
Constant Pressure Calorimetry
By carrying out a
reaction in aqueous
solution in a simple
calorimeter such as this
one, one can indirectly
measure the heat
change for the system
by measuring the heat
change for the water in
the calorimeter.
Thermochemistry
Constant Pressure Calorimetry
Because the specific
heat for water is well
known (4.184 J/mol-K),
we can measure ΔH for
the reaction with this
equation:
Thermochemistry
Using Specific Heat Capacity
A 192-g piece of copper is heated to
100.0 C in boiling water and then
dropped into a beaker containing 751 g
of water at 4.0 C. What is the final
temperature of the copper and water
after thermal equilibrium is reached if
the specific heat of copper is 0.385 J/
gK? (Assume no heat is lost to warm
Thermochemistry
the beaker or the surroundings.)
Another Example
A 15.5-g piece of chromium, heated to
100.0°C, is dropped into 55.5 g of water
at 16.5°C. The final temperature of the
metal and water is 18.9°C. What is the
specific heat capacity of chromium?
(Assume no heat is lost to warm the
container or the surroundings.)
Thermochemistry
Bomb Calorimetry
Reactions can be
carried out in a
sealed _______
such as this one,
and measure the
heat absorbed by
the water.
Thermochemistry
Bomb Calorimetry
•  Because the _________
in the bomb calorimeter is
constant, what is
measured is really the
change in internal energy,
______________
•  For most reactions, the
difference is very _____.
Thermochemistry
Using a Bomb Calorimeter
Methylhydrazine (CH6N2) is commonly used as a liquid
rocket fuel. The combustion of methylhydrazine is as
follows:
2CH6N2 + 5O2 → 2N2 + 2CO2 + 6H2O
When 4.00 g of methylhydrazine is combusted in a
bomb calorimeter, the temperature of the calorimeter
increases from 25.00°C to 39.50°C. In a separate
experiment the heat capacity of the calorimeter is
measured to be 7.794 kJ/°C. What is the heat of
reaction for the combustion of a mole of CH6N2 in this
calorimeter? Was the reaction endo- or exothermic?
Thermochemistry
Hess’s Law
•  ΔH is well known for many reactions,
and it is ___________________ ΔH
for every reaction in which we are
interested.
•  However, we can _______ ΔH using
ΔH values that are _______ and the
properties of enthalpy.
Thermochemistry
Hess’s Law
Hess’s law states that
“If a reaction is carried
out in a series of
steps, ΔH for the
overall reaction
__________________
__________________
________________.”
Thermochemistry
Hess’s Law
Because ΔH is a state
function, the
___________________
depends only on the
______________________
______________________
_____________________.
Thermochemistry
Enthalpies of Formation
An enthalpy of formation, ΔHf, is defined
as the enthalpy change for the reaction
in which a compound is made from its
constituent elements in their elemental
forms.
Thermochemistry
Standard Enthalpies of Formation
Standard enthalpies of formation, ΔHf,° are
measured under standard conditions.
Thermochemistry
Calculation of ΔH
C3H8 (g) + 5 O2 (g) ⎯→ 3 CO2 (g) + 4 H2O (l)
•  Imagine this as occurring
in 3 steps:
C3H8 (g) ⎯→ 3 C(graphite) + 4 H2 (g)
3 C(graphite) + 3 O2 (g) ⎯→ 3 CO2 (g)
4 H2 (g) + 2 O2 (g) ⎯→ 4 H2O (l)
Thermochemistry
Calculation of ΔH
C3H8 (g) + 5 O2 (g) ⎯→ 3 CO2 (g) + 4 H2O (l)
•  Imagine this as occurring
in 3 steps:
C3H8 (g) ⎯→ 3 C(graphite) + 4 H2 (g)
3 C(graphite) + 3 O2 (g) ⎯→ 3 CO2 (g)
4 H2 (g) + 2 O2 (g) ⎯→ 4 H2O (l)
Thermochemistry
Calculation of ΔH
C3H8 (g) + 5 O2 (g) ⎯→ 3 CO2 (g) + 4 H2O (l)
•  Imagine this as occurring
in 3 steps:
C3H8 (g) ⎯→ 3 C(graphite) + 4 H2 (g)
3 C(graphite) + 3 O2 (g) ⎯→ 3 CO2 (g)
4 H2 (g) + 2 O2 (g) ⎯→ 4 H2O (l)
Thermochemistry
Calculation of ΔH
C3H8 (g) + 5 O2 (g) ⎯→ 3 CO2 (g) + 4 H2O (l)
•  The sum of these
equations is:
C3H8 (g) ⎯→ 3 C(graphite) + 4 H2 (g)
3 C(graphite) + 3 O2 (g) ⎯→ 3 CO2 (g)
4 H2 (g) + 2 O2 (g) ⎯→ 4 H2O (l)
C3H8 (g) + 5 O2 (g) ⎯→ 3 CO2 (g) + 4 H2O (l)
Thermochemistry
Calculation of ΔH
We can use Hess’s law in this way:
°
°
where n and m are the stoichiometric
coefficients.
Thermochemistry
Calculation of ΔH
C3H8 (g) + 5 O2 (g) ⎯→ 3 CO2 (g) + 4 H2O (l)
ΔH =
=
=
=
[3(-393.5 kJ) + 4(-285.8 kJ)] - [1(-103.85 kJ) + 5(0 kJ)]
[(-1180.5 kJ) + (-1143.2 kJ)] - [(-103.85 kJ) + (0 kJ)]
(-2323.7 kJ) - (-103.85 kJ)
Notice, that by
-2219.9 kJ
definition,
__________________
__________________
__________________
_________________,
and therefore, the
standard enthalpy Thermochemistry
of
O2 is _____.
Another Example
Calculate ΔH for this reaction
2C (s) + H2 (g) → C2H2 (g)
Given the following reactions and their
respective enthalpy changes
C2H2 (g) + 5/2O2 (g) → 2CO2 (g) + H2O (l) ΔH = -1299.6 kJ
C (s) + O2 (g) → CO2 (g) ΔH = -393.5 kJ
H2 (g) + 1/2O2 (g) → H2O (l) ΔH = -1299.6 kJ
Thermochemistry
You Try This One
Suppose you want to know the enthalpy change
for the formation of methane, CH4, from solid
carbon (as graphite) and hydrogen gas:
C(s) + 2H2 (g) → CH4 (g)
ΔH = ?
The enthalpy change for this reaction cannot be
measured in the laboratory because the
reaction is very slow. We can, however,
measure the enthalpy changes for the
combustion of carbon, hydrogen and
methane.
C (s) + O2 (g) → CO2 (g)
ΔH = -393.5 kJ
H2 (g) + ½ O2 (g) → H2O (l) ΔH = -285.8 kJ
CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l) ΔH = -890.3 kJ
Use these energies to obtain ΔH for the
Thermochemistry
formation of methane.
Energy in Foods
Most of the fuel in the
food we eat comes
from carbohydrates
and fats.
Thermochemistry
Fuels
The vast majority
of the energy
consumed in this
country comes
from fossil fuels.
Thermochemistry
Thermochemistry is Clearly
Important to Our Everyday Lives
Thermochemistry