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Chapter 6
Thermochemistry
許富銀
6.1 Chemical Hand Warmers
Thermochemistry: the study of the
relationships between chemistry and energy
 Hand warmers use the oxidation of iron as
the exothermic reaction:

Nature of Energy



Even though chemistry is the study of matter, energy
affects matter.
Energy is anything that has the capacity to do work.
Work is a force acting over a distance.
–


Energy = work = force × distance
Heat is the flow of energy caused by a difference in
temperature.
Energy can be exchanged between objects through
contact.
–
For example, through collisions
Classification of Energy

Kinetic energy is energy of
motion or energy that is
being transferred.

Thermal energy is the
energy associated with
temperature.
–
Thermal energy is a form of
kinetic energy.
Classification of Energy

Potential energy is energy that is stored in an object,
or energy associated with the composition and
position of the object.
–
Energy stored in the structure of a compound is potential
energy.
FIGURE 6.1 The Different Manifestations
of Energy
Conservation of Energy

The law of conservation of
energy states that energy cannot
be created nor destroyed.

When energy is transferred
between objects, or converted from
one form to another, the total
amount of energy present at the
beginning must be present at the
end.
System and Surroundings


We define the system as the material or process
within which we are studying the energy
changes within
We define the surroundings as everything else with
which the system can exchange energy.
Comparing the Amount of Energy in the
System and Surroundings during Transfer

Conservation of energy
means that the amount of
energy gained or lost by the
system has to be equal to the
amount of energy lost or
gained by the surroundings.
Units of Energy

The amount of kinetic energy an
object has is directly proportional to its mass
and velocity.

The SI unit of energy is therefore kg‧m2/s2 ,
defined as the joule (J),
Units of Energy

A calorie (cal) is the amount of energy needed to raise
the temperature of one gram of water 1 °C.
–
–
kcal = energy needed to raise 1000 g of water 1 °C
food Calories = kcals
The First Law of Thermodynamics:
Law of Conservation of Energy

The first law of thermodynamics is the law of
conservation of energy.
–

This means that the total amount of energy in the universe
is constant.
According to the first law, a device that would
continually produce energy with no energy input,
sometimes known as a perpetual motion machine ,
cannot exist.
Energy Flow and Conservation of
Energy

Conservation of energy requires that the sum of the
energy changes in the system and the surroundings
must be zero.
Energyuniverse = 0 = Energysystem + Energysurroundings
 is the symbol that is used to mean change.

Final amount–initial amount
Internal Energy


The internal energy ( E ) of a system is the
sum of the kinetic and potential energies of all
of the particles that compose the system
The change in the internal energy of a system
only
depends on the amount of energy in the
system at the beginning and end.
– A state function is a mathematical
function whose result only depends on the
initial and final conditions, not on the
process used.
Energy Diagrams

Energy diagram: compares the internal
energy of the reactants and the products:
The reactants become products,
which have lower internal energy.
Therefore, energy is given off by
the reaction and E (that is,
Eproducts - Ereactants ) is negative
• If we define the thermodynamic system as the reactants and
products of the reaction
In this reaction E is positive and energy flows into the system
and out of the surroundings .

In the reaction C(s) + O2(g) →
CO2(g), there will be a net
release of energy into the
surroundings.
– −Ereaction = Esurroundings
Internal Energy
Energy Flow in a Chemical Reaction
C(s), O2(g)
CO2(g)
energy
released
Erxn = ─
Surroundings
System
C + O2 → CO2
In the reaction CO2(g) → C(s)
+ O2(g), there will be an
absorption of energy from the
surroundings into the reaction.
Ereaction = − Esurroundings
Internal Energy

C(s), O2(g)
CO2(g)
Surroundings
System
C + O2 → CO2
energy
absorbed
Erxn = +
Summary: Energy Flow
Surroundings
E ─
System
E +
Surroundings
E +
System
E ─
Energy Exchange

Energy is exchanged between the system and
surroundings through heat and work.
 q = heat (thermal) energy
 w = work energy
 q and w are NOT state functions; their value
depends on the process.
TABLE 6.3 Sign Conventions for q , w ,
and ΔE
FIGURE 6.6 Energy, Work, and Heat
FIGURE 6.6 Energy, Work, and Heat
EXAMPLE 6.1 Internal Energy, Heat,
and Work
The firing of a potato cannon provides a good example of the heat and
work associated with a chemical reaction. A potato is stuffed into a long
cylinder that is capped on one end and open at the other. Some kind of fuel is
introduced under the potato at the capped end—usually through a small
hole—and ignited. The potato then shoots out of the cannon, sometimes
flying hundreds of feet, and the cannon emits heat to the surroundings. If the
burning of the fuel performs 855 J of work on the potato and produces 1422 J
of heat, what is E for the burning of the fuel? (Note: A potato cannon can be
dangerous and should not be constructed without proper training and
experience.)
Sol:
Heat Exchange




Heat is the exchange of thermal energy between a
system and surroundings.
Heat exchange occurs when system and surroundings
have a difference in temperature.
Temperature is the measure of the thermal energy
within a sample of matter.
Heat flows from matter with high temperature to matter
with low temperature until both objects reach the same
temperature.
– Thermal equilibrium
Temperature Changes and Heat
Capacity

When a system absorbs heat ( q ), its
temperature changes by ΔT :

the heat absorbed by a system and its
corresponding temperature change are directly
proportional: q α ΔT .
Units of C are J/°C or J/K
Heat Capacity




the heat capacity ( C ) of a system as the quantity of
heat required to change its temperature by 1 ℃
Heat capacity is an extensive property— it depends on
the amount of matter being heated
specific heat capacity (Cs): the amount of heat required
to raise the temperature of 1 gram of the substance by 1
℃.
molar heat capacity: the amount of heat required to
raise the temperature of 1 mole of a substance by 1 ℃.
Heat Capacity

The specific heat capacity of a substance can
be used to quantify the relationship between
the amount of heat added to a given amount of
the substance and the corresponding
temperature increase
EXAMPLE 6.2 Temperature Changes and
Heat Capacity
Suppose you find a penny (minted before 1982, when
pennies were almost entirely copper) in the snow. How
much heat is absorbed by the penny as it warms from the
temperature of the snow, which is -8.0 ℃, to the
temperature of your body, 37.0 ℃? Assume the penny is
pure copper and has a mass of 3.10 g.
Heat Transfer and Final Temperature



When two objects at different temperatures are
placed in contact, heat flows from the material at the
higher temperature to the material at the lower
temperature.
Heat flows until both materials reach the same final
temperature.
The amount of heat energy lost by the hot material
equals the amount of heat gained by the cold
material.
Thermal Energy Transfer

A block of the metal to the water. metal at 55 °C is
added to water at 25 °C.
–

Thermal energy transfers heat from metal to water
The exact temperature change depends on the
following:
–
–
–
The mass of the metal
The mass of water
Specific heat capacities of the metal and of water
EXAMPLE 6.3 Thermal Energy Transfer

A 32.5 g cube of aluminum initially at 45.8 C is
submerged into 105.3 g of water at 15.4 C. What is
the final temperature of both substances at thermal
equilibrium? (Assume that the aluminum and the
water are thermally isolated from everything else.)
Sol:
Work: Pressure–Volume Work

PV work is work caused by a volume change against an
external pressure.

When gases expand, V is positive, but the system is doing
work on the surroundings, so wgas is negative.

As long as the external pressure is kept constant,
Workgas = External Pressure × Change in Volumegas
w = ─PV.
To convert the units to joules use 101.3 J = 1 atm ∙ L.
EXAMPLE 6.4 Pressure–Volume Work

To inflate a balloon you must do pressure–volume work
on the surroundings. If you inflate a balloon from a
volume of 0.100 L to 1.85 L against an external
pressure of 1.00 atm, how much work is done (in joules)?
Exchanging Energy between System
and Surroundings

Exchange of heat energy
q = mass × specific heat × Temperature

Exchange of work
w = −Pressure × Volume
Bomb Calorimeter


It is used to measure E because it is a constant
volume system.
The heat capacity of the calorimeter is the amount of
heat absorbed by the calorimeter for each degree
rise in temperature and is called the calorimeter
constant. (Ccal, kJ/ºC)
the reaction occurs under conditions of constant volume
EXAMPLE 6.5 Measuring Erxn in a Bomb
Calorimeter

When 1.010 g of sucrose (C12H22O11) undergoes combustion in
a bomb calorimeter, the temperature rises from 24.92 C to
28.33 ℃. Find Erxn 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.)
Sol:
6.6 Enthalpy: The Heat Evolved in a
Chemical Reaction at Constant Pressure


The enthalpy ( H ) of a system : as the sum of its
internal energy and the product of its pressure and
volume:
H = E + PV
H is a state function
The enthalpy change, Δ H, of a reaction is the heat evolved in a
reaction at constant pressure.
Δ Hreaction = q reaction at constant pressure
Usually DH and DE are similar in value; the difference is largest
for reactions that produce or use large quantities of gas.
Endothermic and Exothermic Reactions

Exothermic reaction : when ΔH is negative, heat is
being released by the system.

Endothermic reaction : When ΔH is positive, heat is
being absorbed by the system.
EXAMPLE 6.6 Exothermic and
Endothermic Processes
Identify each process as endothermic or
exothermic and indicate the sign of H .
(a) sweat evaporating from skin
(b) water freezing in a freezer
(c) wood burning in a fire
Molecular View of Exothermic Reactions





For an exothermic reaction, the surrounding’s temperature
rises due to a release of thermal energy by the reaction.
This extra thermal energy comes from the conversion of
some of the chemical potential energy in the reactants into
kinetic energy in the form of heat.
During the course of a reaction, existing bonds are broken
and new bonds are made.
The products of the reaction have less chemical potential
energy than the reactants.
The difference in energy is released as heat.
Molecular View of Endothermic Reactions




In an endothermic reaction, the surrounding’s
temperature drops due to absorption of some of its
thermal energy by the reaction.
During the course of a reaction, existing bonds are
broken and new bonds are made.
The products of the reaction have more chemical
potential energy than the reactants.
To acquire this extra energy, some of the thermal energy
of the surroundings is converted into chemical potential
energy stored in the products.
Stoichiometry Involving H:
Thermochemical Equations


The enthalpy change for a chemical reaction,
abbreviated Hrxn , is also called the enthalpy of
reaction or heat of reaction
The enthalpy change in a chemical reaction is an
extensive property.
–
The more reactants you use, the larger the enthalpy change.
when 1 mol of C3H8 reacts with 5 mol of O2 to form 3 mol of
CO2 and 4 mol of H2O , 2044 kJ of heat is emitted
EXAMPLE 6.7 Stoichiometry Involving
ΔH
Measuring H
Calorimetry at Constant Pressure



Reactions done in aqueous
solution are at constant pressure.
The calorimeter is often nested
foam cups containing the solution.
qreaction = - qsolution
= -(masssolution × Cs, solution × T)
Hreaction = qconstant pressure = qreaction
To get Hreaction per mol, divide by the
number of moles.
EXAMPLE 6.8 Measuring Hrxn in a
Coffee-Cup Calorimeter
Sol:
Relationships Involving Hrxn

When reaction is multiplied by a factor, Hrxn is
multiplied by that factor.
–

Because Hrxn is extensive,
C(s) + O2(g) → CO2(g) H = −393.5 kJ
2 C(s) + 2 O2(g) → 2 CO2(g)
H = 2(−393.5 kJ) =
−787.0 kJ.
If a reaction is reversed, then the sign of H is
changed.
CO2(g) → C(s) + O2(g) H = +393.5 kJ
Relationships Involving Hrxn

If a reaction can be
expressed as a series of
steps, then the Hrxn for
the overall reaction is the
sum of the heats of
reaction for each step.
(Hess’s Law)
EXAMPLE 6.9 Hess’s Law
Sol:
6.9 Determining Enthalpies of Reaction
from Standard Enthalpies of Formation


The standard state is the state of a material at a
defined set of conditions.
– Pure gas at exactly 1 atm pressure
– Pure solid or liquid in its most stable form at
exactly 1 atm pressure and temperature of
interest (Usually 25 °C)
– Substance in a solution with concentration 1 M
The standard enthalpy change, H°, is the
enthalpy change when all reactants and products are
in their standard states.
Standard enthalpy of formation

The standard enthalpy of formation, Hf°, is the
enthalpy change for the reaction forming 1 mole of a
pure compound from its constituent elements.
– The elements must be in their standard states.
– The Hf° for a pure element in its standard state =
0 kJ/mol.
TABLE 6.5 Standard Enthalpies (or
Heats) of Formation, △Hfo , at 298 K
Calculating the Standard Enthalpy
Change for a Reaction

The standard enthalpy of formation corresponds to
the formation of a compound from its constituent
elements in their standard states

The decomposition of a compound into its
constituent elements in their standard states:
Calculating Standard Enthalpy Change
for a Reaction

Any reaction can be written as the sum of formation
reactions (or the reverse of formation reactions) for
the reactants and products.

The H° for the reaction is then the sum of the Hf°
for the component reactions.
H°reaction =  n Hf°(products) −  n Hf°(reactants)
 means sum.
n is the coefficient of the reaction.
EX:
Calculating H° from Tabulated Values
of Hf°
EXAMPLE 6.11 Ho rxn and Standard
Enthalpies of Formation
Sol:
Energy Use and the Environment

Most comes from the combustion of fossil fuels.
–
Combustible materials that originate from ancient life.
C(s) + O2(g) → CO2(g)
H°rxn = −393.5 kJ
CH4(g) +2 O2(g) → CO2(g) + 2 H2O(g)
H°rxn = −802.3 kJ
C8H18(g) +12.5 O2(g) → 8 CO2(g) + 9 H2O(g) H°rxn=−5074.1 kJ


Because of additives and
impurities in the fossil fuel,
incomplete combustion, and side
reactions, harmful materials are
added to the atmosphere when
fossil fuels are burned for energy.
Therefore, fossil fuel emissions
contribute to air pollution, acid rain,
and global warming.
Global Warming



CO2 is a greenhouse gas.
– It allows light from the sun to reach
Earth, but does not allow the heat
(infrared light) reflected off Earth to
escape into outer space.
CO2 levels in the atmosphere have
been steadily increasing.
Current observations suggest that the
average global air temperature has
risen 0.6 °C in the past 100 years.
Renewable Energy

Our greatest unlimited supply of energy is the sun.

New technologies are being developed to capture the
energy of sunlight.
–
–
Parabolic troughs, solar power towers, and dish engines
concentrate the sun’s light to generate electricity.
Solar energy is used to decompose water into H2(g) and O2(g);
the H2 can then be used by fuel cells to generate electricity.
H2(g) + ½ O2(g) → H2O(l)
H°rxn = −285.8 kJ