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Thermochemistry
AP Chemistry Unit 6

Thermodynamics
Thermodynamics is the study of energy and its
interconversions.

Energy
Energy is the essence of our very existence.
 The food we eat gives us the energy to live, work and
play.
 Coal and oil power manufacturing, transportation and
our modern industrialized civilization.

Fossil Fuels
In the past, huge quantities of carbon-based fossil fuels
have been available for the taking.
 This abundance has led to a world society with a
voracious appetite for energy, consuming millions of
barrels of petroleum every day.
 We are now dangerously dependent on the dwindling
supplies of oil. In order to maintain our current
standard of living, energy alternatives must be found.
To find these alternatives, we need to know the
relationship between chemistry and energy.

Chemical Energy
Just as energy is important to our society on a
macroscopic scale, it is critically important to each
living organism on a microscopic scale.
 The living cell is a miniature chemical factory powered
by energy from chemical reactions. The process of
cellular respiration extracts the energy stored in sugars
and other nutrients to drive the various tasks of the cell.
 The conversion of energy from one form to another,
macroscopically and microscopically, are processes
governed by the same principles

Energy
Energy is defined as the capacity to do work or to
produce heat.
 We will concentrate specifically on the heat transfer
that accompanies chemical processes.
 Energy can be classified as either potential or kinetic
energy.

Potential Energy
Potential energy is energy due to position or
composition.
 Example: Water behind a dam has potential energy
that can be converted to work when the water flows
down through turbines, thereby creating electricity.
 Attractive and repulsive forces also lead to potential
energy.

Kinetic Energy
Kinetic energy of an object is energy due to the motion
of the object and depends on the mass of the object
(m) and its velocity (v).
 KE = ½ mv2

Conservation of Energy
The law of conservation of energy states that energy
can be converted from one form to another but can be
neither created nor destroyed.
 Energy of the universe is constant
 Also called the first law of thermodynamics

Conversion of Energy
 A has more potential energy than B
 A transfers energy to B when released
 B ends with more potential energy
 The final energy = initial energy
 transfer of energy to surface
 frictional heating

Conversion of Energy
There are two ways to transfer energy: Work and heat.
 Work is defined as force acting over a distance. Work
is required to raise B from its original position to its final
one. The original energy stored in A is transferred to B
through work.

Heat vs. Temperature
Temperature and heat are very different.
 Temperature is a property that reflects the random
motions of the particles in a particular substance.
 Heat involves the transfer of energy between two
objects due to temperature difference.
 Heat is not a substance contained by an object.

Pathway
The way that the transfer of energy is divided between
work and heat depends on the specific conditions
called the pathway.
 Energy change is independent of the pathway, because
total energy is constant. Work and heat are both
dependent on the pathway; these amounts will differ.

State Function
A state function (or state property) refers to a property of
the system that depends only on its present state.
 The value of a state function does not take into account
how the system arrived at the present state.
 A change in this function (property) in going from one
state to another state is independent of the particular
pathway taken between the two states.
 Energy is a state function, work and heat are not.

State Function Example
 While traveling from Chicago (elevation 674 ft) to
Denver (elevation 5280 ft), the change in elevation is
always 5280 – 674 = 4606 ft. regardless of the route
taken between the two cities. The distance traveled,
depends on how you make the trip. Thus elevation is a
function that does not depend on the route (pathway)
but distance is pathway dependent.
 Elevation is a state function and distance is not.

Chemical Energy
When discussing energy in chemical reactions we divide
everything into two parts:
 The system is the part of the universe on which we
wish to focus attention.
 In chemical reactions we define the system as the
reactants and products of the reaction.
 The surroundings include everything else in the
universe.
 In a chemical reaction, the surroundings consist of the
reaction container, the room and anything else.

Exothermic
When a reaction results in the evolution of heat, it is said
to be exothermic; Energy flows out of the system.
 Heat/Energy is a product.

Endothermic
Reactions that absorb energy from the surroundings are
said to be endothermic; Energy flow is into a system.
 Energy/Heat is a reactant.

Chemical Energy
Where does the energy, released as heat, come from in
an exothermic reaction?
 It is the difference in potential energies between the
products and the reactants.
 Total energy is conserved so the energy gained by the
surroundings must be equal to the energy lost by the
system.

Chemical Energy
This always holds true:
In any exothermic reaction some of the potential energy
stored in the chemical bonds is being converted to
thermal energy (random kinetic energy) via heat.

Chemical Energy
Example: For the combustion of methane ΔPE represents the
change in potential energy stored in the bonds of the
products as compared with the bonds of the reactants. This
is the difference between the energy required to break the
bonds in the reactants and the energy released when the
bonds in the products are formed.

Chemical Energy
 In an exothermic process, the bonds in the products
are stronger than those of the reactants.
 More energy is released by forming the new bonds in
the products than is consumed to break the bonds in
the reactants.
 The quantity ΔPE is transferred to the surroundings
through heat.

Chemical Energy
 In an endothermic reaction, energy that flows into the
system as heat is used to increase the potential energy
of the system.
 The products have higher potential energy (weaker
bonds) than the reactants.

Internal energy
The internal energy (E) of a system is the sum of the
kinetic and potential energies of all the “particles” of the
system.
 The internal energy of a system can be changed by a
flow of work, heat or both.
 ΔE=q+w
 ΔE is the change in the system’s internal energy
 q =heat; w= work

Thermodynamics
 Thermodynamic quantities consist of two parts:
 the magnitude of change
 sign; indicating the direction of the energy flow
 when determining the direction of energy flow, take
the “point of view” of the system.

Endothermic
 ΔE=q+w
 Endothermic process:
 energy flows into the system
 q is +, w is + ; the system’s energy is increasing

Exothermic
 ΔE=q+w
 Exothermic process:
 energy flows out of the system
 q is −, w is −; the system’s energy is decreasing.

Internal Energy Example
Calculate ΔE for a system undergoing an endothermic
process in which 15.6 kJ of heat flows and where 1.4
kJ of work is done on the system.
 ΔE= q + w
 q and w are (+)
 The system gained 17.0 kJ of energy.

Chemical Work Processes
A common type of work associated with chemical
processes is work done by a gas (through expansion)
or work done to a gas (through compression).
 Example: combustion of gasoline in an automobile
engine that causes the motion of the pistons.

Pressure vs. Work
 Pressure is the force acting on the piston of area (A).
(external pressure)
 P = F/A
 Work is a force acting over a distance (h; height)
 w = FΔh
 Change in volume is area x height
 ΔV = A Δh
 Combine equations:
 w = PΔV; work magnitude only.

Work
 The sign of work depends upon the energy flow of the
system.
 w = −PΔV
 If gas is expanding, what is the sign of the system?
 Change in volume is positive.
 Gas expanding is doing work on the surroundings; w
is −

Work
 The sign of work depends upon the energy flow of the
system.
 w = −PΔV
 If a gas is compressed, work is being done on the
system.
 ΔV is negative
 w is positive.

Work Example
Calculate the work associated with the expansion of a gas
from 46 L to 64 L at a constant external pressure of 15
atm.
 w = −PΔV
 P = 15 atm; ΔV = +18L
 w = −270 Latm

Energy, Heat and Work
Example
A balloon is being inflated to its full extent by heating the air
inside it. In the final stages of this process, the volume of
the balloon changes from 4.00 x 106 L to 4.50 x 106 L by
the addition of 1.3 x 108 J of energy as heat. Assuming
that the balloon expands against a constant pressure of
1.0 atm, calculate ΔE for the process. (To convert
between Latm and J, use 1 Latm = 101.3 J.)

Energy, Heat and Work
Example
V1 = 4.00 x 106 L
V2 = 4.50 x 106 L
q = 1.3 x 108 J
P = 1.0 atm
1 Latm = 101.3 J
ΔE = q + w
w = −PΔV
ΔE = 7.9 x 107J

Enthalpy
Enthalpy, H, is defined as H = E + PV
 E is internal energy of the system
 P is the external pressure on the system
 V is the volume of the system
 Enthalpy is the combination of internal energy and the
work done on/by the system.
 Since internal energy, pressure, and volume are all
state functions, enthalpy is also a state function.

Enthalpy
 ΔE = q + w
 ΔE = q − PΔV
 q = ΔE + PΔV
 ΔH = ΔE + PΔV
 ΔH = qp; at a constant pressure (qp)
 At a constant pressure, the flow of heat is a measure of
the change in enthalpy for the system. Therefore, heat
of reaction and change in enthalpy are used
interchangeably.

Enthalpy Change
ΔH = Hproducts – Hreactants
 If ΔH is positive, (products have more enthalpy than
reactants) heat will be absorbed by the system.
 Reaction is endothermic.
 If ΔH is negative, heat is released from the system.
 Reaction is exothermic.

Enthalpy Example
When 1 mole of methane (CH4) is burned at constant
pressure, 890 kJ of energy is released as heat.
Calculate ΔH for a process in which a 5.8 g sample of
methane is burned at constant pressure.
 q = ΔH = −890 kJ/mol CH4
 mass = 5.8 g CH4; molar mass = 16.04 g
 ΔH = −320 kJ

Calorimetry
 Calorimeter – is the device used experimentally to
determine the heat associated with a chemical
reaction.
 Calorimetry is the science of measuring heat.
 based on observing the temperature change when a body
absorbs or discharges energy as heat
 relates energy to heat capacity

Coffee Cup Calorimeter
The measurement of heat using a
simple calorimeter is an example
of constant-pressure
calorimetry. The pressure of
the system, atmospheric
pressure, remains constant.
 used to determine changes in
enthalpy during a chemical
process in solution.

Heat Capacity
Some substances respond differently to being heated.
One substance might require a great deal of heat
energy to raise its temperature by one degree, whereas
another will exhibit the same temperature change after
absorbing relatively little heat. This is also dependent
on the amount of a substance.
 Heat capacity, C, of a substance is a measure of this
property:
 C = heat absorbed/ increase in temperature
 amount of the substance must be specified.

Specific Heat
Specific heat (s) is the energy required to raise the
temperature of one gram of a substance by one degree
Celsius.
 units are J/°Cg or J/Kg
 qp = s x m x ΔT
Molar heat is the energy required to raise the temperature of
one mole of a substance by one degree Celsius
 units are J/°Cmol or J/Kmol

Calorimetry Example
You mix 50.0 ml of 1.0 M HCl at 25.0°C with 50.0 ml of 1.0
M NaOH, also at 25.0°C, in a calorimeter. After the
reactants are mixed by stirring, the temperature is
observed to increase 31.9°C. Specific heat of water is
4.18 J/°Cg. Calculate the heat (change in enthalpy)
for the neutralization reaction.

Calorimetry Example
50.0 ml of 1.0 M HCl at 25.0°C
50.0 ml of 1.0 M NaOHat 25.0°C
Temp increases to 31.9°C.
Specific heat of water is 4.18 J/°Cg; density= 1g/ml
Calculate heat/ΔH.
H+ + OH−  H2O
Change in heat = s  mass ΔT
-2.9 x 103 J

Extensive vs. Intensive
 Change in heat or enthalpy is an extensive property.
 It is directly related to the amount of a substance that you
have.
 Intensive properties are not related to the amount of a
substance.
 Temperature

Calorimetry Example
When 1.00 L of 1.00M Ba(NO3)2 solution at 25.0°C is
mixed with 1.00L of 1.00 M Na2SO4 solution at 25.0°C
in a calorimeter, the white solid BaSO4 forms, and the
temperature of the mixture increases to 28.1°C.
Assuming that the calorimeter absorbs only a negligible
quantity of heat, that the specific heat of the solution is
4.18 J/°Cg, and the density of the final solution is 1.0
g/ml, calculate the enthalpy change per mole of BaSO4
formed.

Calorimetry Example
1.00 L of 1.00M Ba(NO3)2 at 25.0°C
1.00L of 1.00 M Na2SO4 at 25.0°C
BaSO4 forms, and temperature increases to 28.1°C.
s = 4.18 J/°Cg; d=1.0 g/ml,
Calculate the ΔH/mol Ba2+
(aq) + SO4
2
(aq)
-  BaSO4(s)
qp = −26 kJ
ΔH = −26 kJ/mol

Constant Volume Calorimetry
When the volume remains constant during a reaction, (in
a closed container), a bomb calorimeter is used.
 Constant-Volume Calorimetry often involves the heat
capacity of the “bomb” calorimeter as a total system.
 Since w = −PΔV; if change in volume is 0, work would
also be 0.
 ΔE = q + w = qv

Constant Volume Calorimetry
Example
Calculate the energy of combustion of octane (C8H18), a
component of gasoline, if a 0.5269g sample of octane is
placed in a bomb calorimeter that has a heat capacity of
11.3 kJ/°C. The octane is ignited in the presence of
excess oxygen, and the temperature increase of the
calorimeter is 2.25°C.
 Energy released = Heat capacity x ΔT = 25.4 kJ
 Moles of octane = .5269g = 4.614 x 10-3 mol
 Energy released / mol = 5.50 x 103 kJ/mol
 reaction is exothermic; sign is negative
 q = -5.50 x 103 kJ/mol

Constant Volume Example
It has been suggested that hydrogen gas obtained by the
decomposition of water might be a substitute for natural
gas (principally methane). To compare the energies of
combustion of these fuels, the following experiment was
carried out using a bomb calorimeter with a heat capacity
of 11.3 kJ/°C. When a 1.50 g sample of methane gas
was burned with excess oxygen in the calorimeter, the
temperature increased by 7.3°C. When a 1.15 g sample
of hydrogen gas was burned with excess oxygen, the
temperature increase was 14.3°C. Compare the
energies of combustion (per gram) for hydrogen and
methane.

Constant Volume Example
bomb heat capacity = 11.3 kJ/°C
1.50 g methane gas; ΔT = 7.3°C
1.15 g hydrogen gas; ΔT = 14.3°C.
Compare the energies of combustion (per gram).
Energy of 1.5g CH4 = 11.3kJ/°C x 7.3°C = 83 kJ
Energy of 1.0 g CH4 = 55 kJ/g
Energy of 1.15g H2 = 11.3 kJ/°C x 14.3°C = 162 kJ
Energy of 1.0 g H2 = 141 kJ/g

Hess’s Law
Since enthalpy is a state function, the change in enthalpy
in going from some initial state to some final state is
independent of the pathway. This means that in going
from a particular set of reactants to a particular set of
products, the change in enthalpy is the same whether
the reaction takes place in one step or in a series of
steps. This is known as Hess’s Law.

Hess’s Example
N2(g) + 2O2(g)  2NO2(g)
 Change in enthalpy, ΔH 68 kJ
N2(g) + O2(g)  2NO(g) ΔH = 180 kJ
2NO(g) + O2(g)  2NO2(g) ΔH = −112 kJ
N2(g) + 2O2(g)  2NO2(g) ΔH = 68 kJ

Rules for Hess’s Law
Calculations
 If the reaction is reversed, the sign of ΔH is also
reversed.
 N2(g) + O2(g)  2NO(g) ΔH = 180 kJ
 2NO(g)  N2(g) + O2(g) ΔH = −180 kJ
 The magnitude of ΔH is directly proportional to the
quantities of reactants and products in a reaction. If the
coefficients in a balanced reaction are multiplied by an
integer, the value of ΔH is multiplied by the same
integer.
 2N2(g) + 2O2(g)  4NO(g) ΔH = 360 kJ

Hess’s Law Example
Two forms of carbon are graphite, the soft,
black, slippery material used in “lead”
pencils and as a lubricant for locks, and
locks, and diamond, the brilliant, hard
gemstone. Using enthalpies of combustion
for graphite (-394 kJ/mol) and diamond (-
396 kJ/mol), calculate ΔH for the conversion
of graphite to diamond:
Cgraphite(s)  Cdiamond(s)

Cgraphite(s)  Cdiamond(s)
Cgraphite(s) + O2  CO2(g) ΔH = −394 kJ
Cdiamond(s) + O2  CO2(g) ΔH = −396 kJ
ΔH = 2 kJ

Diborane (B2H6) is a highly reactive boron hydride that was
once considered as a possible rocket fuel for the U.S space
program. CalculateΔH for the synthesis of diborane from
its elements, according to the equation
2B(s) + 3H2(g)  B2H6(g)
Use the following data:

2B(s) + 3H2(g)  B2H6(g)
2B(s) + O2(g)  B2O3(s) -1273
kJ
B2H6(g) + 3O2(g)  B2O3(s) + 3H2O(g) -2035
kJ
H2(g) + ½O2(g)  H2O(l) -286 kJ
H2O(l)  H2O(g) 44 kJ
3
2
ΔH = 36 kJ/mol

Standard Enthalpies of
Formation
6.4

Formation
The standard enthalpy of formation (ΔHf°) of a
compound is defined as the change in enthalpy that
accompanies the formation of one mole of a compound
from its elements with all substances in their standard
states.
 A degree symbol on a thermodynamic function, for
example, ΔH°, indicates that the corresponding process
has been carried out under standard conditions. The
standard state for a substance is a precisely defined
reference state.

Formation
 Since thermodynamic functions often depend on the
concentrations (or pressures) of the substances
involved, we must use a common reference state to
properly compare the thermodynamic properties of two
substances. This is especially important because, for
most thermodynamic properties, we can measure only
changes in the property.

Standard State Definitions
 The standard state of a gaseous substance is a
pressure of exactly 1 atm.
 For a pure substance in a condensed state (liquid or
solid), the standard state is the pure liquid or solid.
 For a substance present in a solution, the standard
state is a concentration of exactly 1 M.
 For an element the standard state is the form in which
the element exists under conditions of 1atm and 25°C.

Enthalpy of Formation
 Is stated such that the elements are in their standard
states.
 They are always written so that 1 mole of product is
formed in its standard state.
 ½NO(g) + O2(g)  NO2(g) ΔHf = 34 kJ/mol
 C(s) + 2H2(g) + ½O2(g)  CH3OH(l) ΔHf = −239
kJ/mol

Example of Enthalpy
Calculations
What is the enthalpy of the following reaction (ΔH°reaction):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
−891 kJ

Enthalpy Formation
The process of a chemical reaction that includes
enthalpies of formation can be summarized as follows:
 The enthalpy change for a given reaction can be
calculated by subtracting the enthalpies of formation of
the reactants from the enthalpies of formation of the
products. Remember to multiply the enthalpies of
formation by integers as required by the balanced
equation.
 ΔH°reaction = ΣnpΔHf°(products) − ΣnrΔHf°(reactants)

Enthalpy Problem 1
Using the standard enthalpies of formation listed in table
6.2, calculate the standard enthalpy change for the
overall reaction that occurs when ammonia is burned in
air to form nitrogen dioxide and water. this is the first
step in the manufacture of nitric acid.
4NH3(g) + 7O2(g)  4NO2(g) + 6H2O(l)

Enthalpy Problem 1
4NH3(g) + 7O2(g)  4NO2(g) + 6H2O(l)
−1396 kJ

Enthalpy Problem 2
Using enthalpies of formation, calculate the standard
change in enthalpy for the thermite reaction:
2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)
This reaction occurs when a mixture of powdered
aluminum and iron (III) oxide is ignited with a
magnesium fuse.

Enthalpy Problem 2
2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)
−850. kJ

Enthalpy Problem 3
Until recently, methanol (CH3OH) was used as a fuel in
high-performance engines in race cars. Using the data
in Table 6.2, compare the standard enthalpy of
combustion per gram of methanol with that per gram of
gasoline. Gasoline is actually a mixture of compounds,
but assume for this problem that gasoline is pure liquid
octane (C8H18)

Enthalpy Problem 3
Methanol: 2CH3OH(l) + 3O2  2CO2(g) + 4H2O(l)
−22.6 kJ/g

Enthalpy Problem 3
Octane: 2C8H18(l) + 25O2  16CO2(g) + 18H2O(l)
−47.7 kJ/g

Enthalpy Problem 3
 The enthalpy of combustion per gram of octane is
approximately twice that per gram of methanol. On this
basis, gasoline appears to be superior to methanol for
use in a racing car.

Present Energy Sources
Woody plants, coal, petroleum, and natural gas hold a
vast amount of energy that originally came from the
sun. By the process of photosynthesis, plants store
energy that can be claimed by burning the plants
themselves or the decay products that have been
converted over millions of years to fossil fuels.
Although the United States currently depends heavily
on petroleum for energy, this dependence is a relatively
recent phenomenon.

Petroleum and Natural Gas
Petroleum and natural gas were most likely formed from
the remains of marine organisms that lived
approximately 500 million years ago. Petroleum is a
thick, dark liquid composed mostly of compounds
called hydrocarbons that contain carbon and hydrogen.
Carbon is unique among elements in the extent to
which it can bond to itself to form chains of various
lengths.

Natural gas usually associated with petroleum deposits,
consists mostly of methane, but it always contains
significant amounts of ethane, propane and butane.
The composition of petroleum varies from 5 to 25 or more
Carbons in the chains. For petroleum to be used
efficiently, the petroleum must be separated into
fractions by boiling. The lighter molecules, those with
lower boiling points, can be boiled off.

The petroleum era began when the demand for lamp oil
during the Industrial Revolution outstripped the
traditional sources: animal fats and whale oil. In
response to this increased demand, Edwin Drake
drilled the first oil well in 1859 at Titusville,
Pennsylvania. The petroleum from this well was
refined to produce kerosene (fraction C10 – C18), which
served as an excellent lamp oil.

Gasoline (fraction C5 – C10) had limited used and was often
discarded. However, this situation soon changed. The
development of the electric light decreased the need for
kerosene, and the advent of the “horseless carriage” with its
gasoline-powered engine signaled the birth of the gasoline age.
As gasoline became more important, new ways were sought to
increase the yield of gasoline obtained from each barrel or
petroleum. William Burton invented a process at Standard Oil of
Indiana called pyrolytic (high-temperature) cracking. In this
process, the heavier molecules of the kerosene fraction are
heated to about 700°C, causing them to break (crack) into
smaller molecules of hydrocarbons in the gasoline fraction.

Because of the uneven burning of the gasoline then
available, these engines “knocked,” producing unwanted
noise and even engine damage. The addition of
tetraethyl lead, an “anti-knock” agent, to gasoline
became a common practice, and by 1960, gasoline
contained as much as 3 g of lead per gallon. As we
have discovered so often in recent years, technological
advances can produce environmental problems. To
prevent air pollution from automobile exhaust, catalytic
converters have been added to car exhaust systems.
The effectiveness of these converters, however, is
destroyed by lead.

The use of leaded gasoline also greatly increased the
amount of lead in the environment, where it can be
ingested by animals and humans. For these reasons,
the use of lead in gasoline has been phased out,
requiring extensive (and expensive) modifications of
engines and of the gasoline refining process.

Coal
Coal was formed from the remains of plants that were
buried and subjected to high pressure and heat over
long periods of time. Plant materials have a high
content of cellulose, a complex molecule whose
empirical formula is CH2O but whose molar mass is
around 500,000 g/mol. After the plants and trees that
flourished on the earth at various times and places died
and were buried, chemical changes gradually lowered
the oxygen and hydrogen content of the cellulose
molecules.

Coal
Coal “matures” through four stages: lignite,
subbituminous, bituminous, and anthracite. Each stage
has a higher carbon to oxygen and carbon to hydrogen
ratio; that is the relative carbon content gradually
increases. The energy available from the combustion
of a given mass of coal increases as the carbon
content increases. Therefore, anthracite is the most
valuable coal, and lignite the least valuable.

Coal
Coal is an important and plentiful fuel in the US, currently
furnishing approximately 23 % of our energy. As the supply of
petroleum dwindles, the share of the energy supply from coal
is expected to increase. However, coal is expensive and
dangerous to mine underground, and the strip mining of fertile
farmland in the Midwest or of the scenic land in the West cause
obvious problems. In addition, the burning of coal, especially
high sulfur coal, yields air pollutants such as sulfur dioxide,
which in turn can lead to acid rain. However, even if coal were
pure carbon, the carbon dioxide produced when it was burned
could still have significant effect of the earth’s climate.

CO2 and Climate
The earth receives a tremendous quantity of radiant energy
from the sun, about 30% of which is reflected back into
space by the earth’s atmosphere. The remaining energy
passes through the atmosphere to the earth’s surface.
Some of this energy is absorbed by plants for
photosynthesis and some by the oceans to evaporate
water, but most of it is absorbed by soil, rocks, and water,
increasing the temperature of the earth’s surface. This
energy is in turn radiated from the heated surface mainly
as infrared radiation, often called heat radiation.

CO2 and Climate
The atmosphere, like window glass, is transparent to visible
light but does not allow all the infrared radiation to pass
back into space. Molecules in the atmosphere,
principally H2O and CO2, strongly absorb infrared
radiation and radiate it back toward the earth, so a net
amount of thermal energy is retained by the earth’s
atmosphere. In a way, the atmosphere acts like the glass
of a greenhouse, which is transparent to visible light but
absorbs infrared radiation, thus raising the temperature
inside the building. This greenhouse effect is seen
even more spectacularly on Venus, where the dense
atmosphere is thought to be responsible for the high
surface temperature of that planet.

CO2 and Climate
Thus the temperature of the earth’s surface is controlled
to a significant extent by the carbon dioxide and water
content of the atmosphere. The effect of atmospheric
moisture (humidity) is apparent in the Midwest. In
summer, when the humidity is high, the heat from the
sun is retained well into the night, giving very high
nighttime temperatures. On the other hand, in winter,
the coldest temperatures always occur on clear nights,
when the low humidity allows efficient radiation of
energy back into space.

CO2 and Climate
The atmosphere’s water content is controlled by the water cycle
(evaporation and precipitation), and the average remains
constant over the years. However, as fossil fuels have been
used more extensively, the carbon dioxide concentration has
increased by about 16% from 1880 to 1980. Comparisons of
satellite data have now produced evidence that the greenhouse
effect has significantly warmed the earth’s atmosphere. The
data compare the same areas in both 1979 and 1997. The
analysis shows that more infrared radiation was blocked by
CO2, methane, and other greenhouse gases. This could
increase the earth’s average temperature by as much as 3°C,
causing dramatic changes in climate and greatly affecting the
growth of food crops.

CO2 and Climate
How well can we predict long-term effects? Because
weather has been studied for a period of time that is
minuscule compared with the age of the earth, the
factors that control the earth’s climate in the long range
are not clearly understood. For example, we do not
understand what causes the earth’s periodic ice ages.
So it is difficult to estimate the impact of the increasing
carbon dioxide levels.

CO2 and Climate
Although the exact relationship between the carbon
dioxide concentration in the atmosphere and the earth’s
temperature is not known at present, on thing is certain.
The increase in the atmospheric concentration of
carbon dioxide is quite dramatic. We must consider the
implications of this increase as we consider our future
energy needs.

New Energy Sources
As we search for the energy sources of the future, we
need to consider economic, climatic, and supply
factors. There are several potential energy sources:
the sun (solar), nuclear processes (fission and fusion),
biomass (plants), and synthetic fuels. Direct use of the
sun’s radiant energy to heat our homes and run our
factories and transportation systems seems a sensible
long-term goal. But what do we do now? Conservation
of fossil fuels is one obvious step, but substitutes for
fossil fuels also must be found.

Coal Conversion
One alternative energy sources involves using a
traditional fuel – coal – in new ways. Since
transportation costs for solid coal are high, more
energy-efficient fuels are being developed from coal.
One possibility is to produce a gaseous fuel.
Substances like coal that contain large molecules have
high boiling points and tend to be solids or thick liquids.
To convert coal from a solid to a gas therefore requires
reducing the size of the molecules; the coal structure
must be broken down in a process called coal
gasification.

Coal Conversion
Coal gasification is done by treating the coal with oxygen
and steam at high temperatures to break many of the
carbon-carbon bonds. These bonds are replaced by
carbon-hydrogen and carbon-oxygen bonds as the coal
fragments react with the water and oxygen. The
desired product is a mixture of carbon monoxide and
hydrogen called synthetic gas, or syngas, and
methane (CH4) gas. Since all the components of this
product can react with oxygen to release heat in a
combustion reaction, this gas is a useful fuel.

Coal Conversion
One of the most important consideration in designing an
industrial process is efficient use of energy. In coal
gasification, some of the reactions are exothermic.
Other gasification reactions are endothermic, but if
such conditions as the rate of feed of coal, air, and
steam are carefully controlled, the correct temperature
can be maintained in the process without using any
external energy source. That is, an energy balance is
maintained.

Coal Conversion
Although syngas can be used directly as a fuel, it is also
important as a raw material to produce other fuels. For
example, syngas can be converted directly to
methanol. Methanol is used in the production of
synthetic fibers and plastics and also can be used as a
fuel. In addition, it can be converted directly to gasoline.
Approximately half of South Africa’s gasoline supply
comes from methanol produced from syngas.

Coal Conversion
In addition to coal gasification, the formation of coal
slurries is another new use of coal. A slurry is a
suspension of fine particles in a liquid, and coal must
be pulverized and mixed with water to form a slurry.
The slurry can be handled, store, and burned in ways
similar to those used for residual oil, a heavy fuel oil
from petroleum accounting for almost 15% of US
petroleum imports. One hope is that coal slurries might
replace solid coal and residual oil as fuels for
electricity-generating power plants. However, the water
needed for slurries might place an unacceptable
burden on water resources, especially in the western
US.

Hydrogen as a Fuel
Hydrogen has a real advantage over fossil fuels in that
the only product of hydrogen combustion is water; fossil
fuels also produce carbon dioxide. However, even
though it appears that hydrogen is a very logical choice
as a major fuel for the future, there are three main
problems: the cost of production, storage, and
transport.

Hydrogen as a Fuel
 The Main Engines in the
Space Shuttle Endeavour
Use the Hydrogen and
Oxygen as Fuel.

Hydrogen as a Fuel
First let’s look at the production problem. Although
hydrogen is very abundant on earth, virtually none of it
exists as the free gas. Currently, the main source of
hydrogen gas is from the treatment of natural gas with
steam. The reaction is highly endothermic; treating
methane with steam is not an efficient way to obtain
hydrogen for fuel. It would be much more economical
to burn the methane directly.

Hydrogen as a Fuel
A virtually inexhaustible supply of hydrogen exists in the
waters of the world’s oceans. However, the reaction
requires 286 kJ of energy per mole of liquid water, and
under current circumstances, large scale production of
hydrogen from water is not economically feasible.
However, several methods for such production are
currently being studied: electrolysis of water, thermal
decomposition of water, thermochemical decomposition
of water and biological decomposition.

Hydrogen as a Fuel
Electrolysis of water involves passing an electric current
through it. The present cost of electricity makes the
hydrogen produced by electrolysis too expensive to be
competitive as a fuel. However, if in the future we
develop more efficient sources of electricity, this
situation could change.

Hydrogen as a Fuel
Recent research at the University of Minnesota by Lanny
Schimidt and his coworkers suggests that corn could
be a feasible source of hydrogen. In this process the
starch from the corn is fermented to produce alcohol,
which is then decomposed in a special reactor at
140°C with a rhodium and cerium oxide catalyst to give
hydrogen. These scientists indicate that enough
hydrogen gas can be obtained from a few ounces of
ethanol to generate electricity to run six 60 watt bulbs
for an hour.

Hydrogen as a Fuel
Thermal decomposition is another method for producing
hydrogen from water. This involves heating the water
to several thousand degrees, where is spontaneously
decomposes into hydrogen and oxygen. However,
attaining temperatures in this range would be very
expensive even if a practical heat source and suitable
reaction container were available.

Hydrogen as a Fuel
But what about organisms that decompose water without
the aid of electricity or high temperatures? In the
process of photosynthesis, green plants absorb carbon
dioxide and water and use them along with energy from
the sun to produce the substances needed for growth.
Scientists have studied photosynthesis for years,
hoping to get answers to humanity’s food and energy
shortages. At present, much of this research involves
attempts to modify the photosynthetic process so that
plants will release hydrogen gas from water instead of
using the hydrogen to produce complex compounds.

Hydrogen as a Fuel
Small-scale experiments have shown that under certain
conditions plants do produce hydrogen gas, but the
yields are far from being commercially useful. At this
point the economical production of hydrogen remains
unrealized.

Hydrogen as a Fuel
The storage and transportation of hydrogen also present
problems. First, on metal surfaces the H2 molecule
decomposes to atoms. Since the atoms are so small,
they can migrate into the metal, causing structural
changes that make it brittle. This might lead to a
pipeline failure if hydrogen were pumped under high
pressure. An additional problem is the relatively small
amount of energy that is available per unit volume of
hydrogen gas. Although the energy available per gram
of hydrogen is significantly greater than that per gram of
methane, the energy available per given volume of
hydrogen is about one-third that available from the same
volume of methane.

Hydrogen as a Fuel
Although the use of hydrogen as a fuel solves some of the
problems associated with fossil fuels, it does present
some potential environmental problems of its own.
Studies by John M. Eiler and his colleagues at
California Institute of Technology indicate that, if
hydrogen becomes a major source of energy,
accidental leakage of the gas into the atmosphere
could pose a threat. The Cal Tech scientists calculate
that leakage could raise the concentration of H2 in the
atmosphere from its natural level of .5 ppm to more
than 2 ppm.

Hydrogen as a Fuel
As some of the H2 eventually finds its way into the upper
atmosphere, it would react with O2 to form water, which would
increase the number of ice crystals. This could lead to the
destruction of some of the protective ozone because many of
the chemical reactions that destroy ozone occur on the
surfaces of ice crystals. However, as is the usual case with
environmental issues, the situation is complicated. The
scenario suggested by Eiler’s team may not happen because
the leaked H2 could be consumed by soil microbes that use
hydrogen as a nutrient. In fact, Eiler’s studies show that 90%
of the H2 emitted into the atmosphere today from sources such
as motor vehicles and forest fires is eventually absorbed by
soil organisms.

Hydrogen as a Fuel
Could hydrogen be considered as a potential fuel for
automobiles? This is an intriguing question. The
internal combustion engines in automobiles can be
easily adapted to burn hydrogen. In fact, BMW is now
experimenting with a fleet of cars powered by
hydrogen-burning internal combustion engines.
However, the primary difficulty is the storage of enough
hydrogen to give an automobile reasonable range. The
automobile would need a huge tank to hold enough
hydrogen gas to have a typical mileage range.

Hydrogen as a Fuel
Clearly, hydrogen must be stored as a liquid or in some
other way. Is this feasible? Because of its very low
boiling point (20K), storage of liquid hydrogen requires
a superinsulated container that can withstand high
pressures. Storage in this manner would be both
expensive and hazardous because of the potential for
explosion. Thus storage of hydrogen in the individual
automobile as a liquid does not seem practical.

Hydrogen as a Fuel
A much better alternative seems to be the use of metals
that absorb hydrogen to form solid metal hydrides:
H2(g) + M(s)  MH2(s)
To use this method of storage, hydrogen gas would be
pumped into a tank containing the solid metal in
powdered form, where it would be absorbed to form the
hydride, whose volume would be little more than that of
the metal alone. This hydrogen would then be
available for combustion in the engine by release of
H2(g) from the hydride as needed.

Other Energy Alternatives
Many other energy sources are being considered for future use.
The western states, especially Colorado, contain huge deposits
of oil shale, which consists of a complex carbon-based material
called kerogen contained in porous rock formations. These
deposits have the potential of being a larger energy source than
the vast petroleum deposits of the Middle East. The main
problem with oil shale is that the trapped fuel is not fluid and
cannot be pumped. To recover the fuel, the rock must be heated
to a temperature of 250°C or higher to decompose the kerogen
to smaller molecules that produce gaseous and liquid products.
This process is expensive and yields large quantites of waste
rock, which have a negative environmental impact.

Ethanol (C2H5OH) is another fuel with the potential to
supplement, if not replace, gasoline. The most common
method of producing ethanol is fermentation, a process in
which sugar is changed to alcohol by the action of yeasts.
The sugar can come from virtually any source, including
fruits and grains, although fuel-grade ethanol would
probably come mostly from corn. Car engines can burn
pure alcohol or gasohol, an alcohol-gasoline mixture (10%
ethanol in gasoline), with little modification. Gasohol is
now widely available in the US. A fuel called E85, which is
85% ethanol and 15% gasoline, is also widely available for
cars with “flex-fuel” engines.

Methanol (CH3OH), an alcohol similar to ethanol, which
was used successfully for many years in race cars, is
now being evaluated as a motor fuel in California. A
major gasoline retailer has agreed to install pumps at
25 locations to dispense a fuel that is 85% methanol
and 15% gasoline for use in specially prepared
automobiles. The California Energy Commission feels
that methanol has great potential for providing a
secure, long term energy supply that would alleviate air
quality problems.

This Promotion Bus Both Advertises Biodiesel
and Demonstrates its Usefulness

Another potential source of liquid fuels is oil squeezed from
seeds (seed oil). For example, some farmers in North
Dakota, South Africa, and Australia are now using
sunflower oil to replace diesel fuel. Oil seeds, found in a
wide variety of plants, can be processed to produce an oil
composed mainly of carbon and hydrogen, which of
course reacts with oxygen to produce carbon dioxide,
water and heat. It is hoped that oil-seed plants can be
developed that will thrive under soil and climatic
conditions unsuitable for corn and wheat. The main
advantage of seed oil as a fuel is that it is renewable.
Ideally, fuel would be grown just like food crops.

Spontaneous Process
A process is said to be spontaneous if it occurs
without outside intervention. This process can be fast
or slow.
 Examples: If exposed to air and moisture, steel
rusts spontaneously. However, the iron oxide in
rust does not spontaneously change back to iron
metal and oxygen gas.

Spontaneous Process
It was once believed that a reaction would be
spontaneous if it were exothermic.
 Energy release is an important factor in
spontaneity, but not the total answer.
 Example: Ice melts spontaneously at
temperatures greater than 0°C, but is an
endothermic process.

Entropy
Entropy, S, is the driving force for a spontaneous
process. Entropy can be viewed as a measure of
molecular randomness or disorder.
 The natural progression of things is from order
to disorder. From lower entropy to higher
entropy.
 Example: Your room! There are many more
ways for things to be out of place than for them
to be in their place.

Entropy
Entropy is a thermodynamic function that
describes the number of arrangements (positions
and/or energy levels) that are available to a
system existing in a given state.
 Entropy is closely associated with probability.
 Nature spontaneously proceeds toward the
states that have the highest probabilities of
existing.

Entropy Example
How many ways can each state
be achieved?
 Arrangements I and V can be
achieved in only one wary – all
molecules must be in one end
(one microstate).
 Arrangements II and IV can be
achieved in four ways. (four
microstates).

Entropy Example
How many ways can each state
be achieved?
 Arrangement III can be
achieved in six ways (six
microstates).
 Which arrangement is most
likely to occur?
 Arrangement III is most
probable with six microstates.

Entropy Example
 Each molecule has a 1:2
chance of being on the left
side. Two molecules being on
the left have a 1:4 chance.
Eventually, with larger number
of molecules, the probability
of finding all molecules on one
side would “never” occur.

Entropy Example
 If something is highly
improbable; this process does
not occur spontaneously.
 Gas expands to a uniform
distribution because of
positional probability. The
expanded state has the
highest positional probability
and the largest entropy.

Changes of State
Changes of state illustrate positional probability.
In general, positional entropy increases in going
from solid to liquid to gas.
 In a solid state the molecules are close
together, with relatively few positions available
to them; in a gas state the molecules are far
apart, with many more positions available.

Changes of State
Changes of state illustrate positional probability.
In general, positional entropy increases in going
from solid to liquid to gas.

Mixing Solutions
Entropy change associated with the mixing of two
pure substances is expected to be positive.
 When two liquids are mixed to form a solution,
the molecules of each liquid have more
available volume and thus more available
position.

Example Problem
 For each of the following pairs, choose the
substance with the higher positional entropy
(per mole) at a given temperature:
 Solid CO2 or gaseous CO2
 N2 gas at 1 atm or N2 gas at 1.0 x 10-2 atm

Example Problem
Predict the sign of the entropy change for each of
the following processes:
 Solid sugar is added to water to form a
solution. ΔS+
 Iodine vapor condenses on a cold surface to
form crystals ΔS−

Entropy and 2nd Law of
Thermodynamics
6.8

2nd Law of Thermodynamics
In any spontaneous process there is always an
increase in the entropy of the universe.
 Energy is conserved in the universe, but
entropy is not.
 Entropy of the universe is increasing.

2nd Law of Thermodynamics
ΔSuniv = ΔSsystem + ΔSsurroundings
 If ΔSuniv is positive, the entropy of the universe
increases, and the process is spontaneous.
 If ΔSuniv is negative, the process is
spontaneous in the opposite direction.
 If ΔSuniv is zero, the process has no tendency to
occur, and the system is at equilibrium.

Temperature and Spontaneity
An exothermic process transfers energy as heat
to the surroundings, where it becomes
thermal/kinetic energy associated with random
motions of atoms. Energy flow into the
surroundings increases entropy in the
surroundings.
 ΔSsurr is positive.

An endothermic process allows heat to flow from
the surroundings into the system and the random
motions of the atoms in the surroundings
decrease, decreasing entropy in the
surroundings.
 ΔSsurr is negative.

When ΔSsys and ΔSsurr are in opposition (opposite
signs) the ΔSuniv is dependent on temperature.
 Example: Vaporization of water is endothermic:
ΔSsys is +, ΔSsurr is −
 At 1 atm, water changes spontaneously to a
gas above 100°C; below 100°C, condensation
is spontaneous.

Entropy changes in the surroundings are
primarily determined by heat flow
 The impact of the transfer of a given quantity of
energy as heat to or from the surroundings will
be greater at lower temperatures.

Summary
1. The sign of ΔSsurr depends on the direction of
heat flow:
1. Exothermic = ΔSsurr is +
2. Endothermic = ΔSsurr is −
2. The magnitude of ΔSsurr depends on the
temperature.
1. The tendency for the system to lower its
energy becomes a more important driving
force at lower temperatures.

Summary
ΔSsurr =
The minus sign is necessary because the sign of
ΔH is determined with respect to the reaction
system, and this equation expresses a property
of the surroundings.
For example, if the reaction is exothermic, ΔH
has a negative sign, but since heat flows into the
surroundings, ΔSsurr is positive.

Example
In the metallurgy of antimony, the pure metal is recovered via
different reactions, depending on the composition of the ore.
For example, iron is used to reduce antimony in sulfide ores:
 Sb2S3(s) + 3Fe(s)  2Sb(s) + 3FeS(s) ΔH = -125kJ
Carbon is used as the reducing agent for oxide ores:
 Sb4O6(s) + 6C(s)  4Sb(s) + 6CO(g) ΔH = 778kJ
Calculate ΔSsurr for each of these reactions at 25°C and 1atm.

Example
 Sb2S3(s) + 3Fe(s)  2Sb(s) + 3FeS(s) ΔH = -125kJ
ΔSsurr =419 J/K

Example
 Sb4O6(s) + 6C(s)  4Sb(s) + 6CO(g) ΔH = 778kJ
ΔSsurr = − 2.61 x 103 J/K

Free Energy
Gibbs Free Energy represents the energy difference
between enthalpy and the temperature and entropy of
the surroundings.
 ΔG = ΔH−TΔS
 H is enthalpy, T is Kelvin temperature, and S is entropy
 All quantities refer to the system.
 A process (at constant T and P) will be spontaneous if
ΔG is negative.

Gibbs Example
H2O(s)  H2O(l) ΔH° = 6.03 x 103 J/mol
ΔS° (ΔSsys) = 22.1 J/Kmol
Is this reaction spontaneous at 10°C?
Is this reaction spontaneous at −10°C?

Gibbs Example 2
At what temperatures is the following process
spontaneous at 1 atm?
Br2(l)  Br2(g)
 ΔH° = 31.0 kJ/mol and ΔS° = 93.0 J/Kmol
 In other words, what is the normal boiling point of liquid
Br2?
 Hint: Set G=0
333K

Entropy Changes in Chemical
Reactions
17.5

Entropy in Reactions
Entropy Summary:
 The entropy changes in the surroundings are
determined by the heat flow that occurs as the reaction
takes place.
 The entropy changes in the system are determined by
positional probability.

Example
N2(g) + 3H2(g)  2NH3(g)
 Fewer molecules mean fewer possible configurations.
Entropy for the system is less.
 In general, when a reaction involves gaseous
molecules, the change in positional entropy is
dominated by the relative numbers of molecules of
gaseous reactants and products.

Example Problem
Predict the sign of ΔS° for each of the following reactions.
 CaCO3(s)  CaO(s) + CO2(g)
 +
 2SO2(g) + O2(g)  2SO3(g)
 −

3rd Law of Thermodynamics
Values for entropy are assigned by comparing individual
substances to 0 entropy.
 A solid with perfect crystal structure at 0 Kelvin is
considered 0 entropy because at this state molecular
motion ceases.
The third law of thermodynamics states that the entropy of
a perfect crystal at 0 K is zero.
 As the temperature increases, the random vibrational
motions increases and disorder increases.

Entropy
Entropy is a state function of the system and can be
calculated by using stand entropy values from the
Appendix.
ΔS°reaction = ΣnpS°products − ΣnpS°reactants

Example Problem
Calculate ΔS° at 25°C for the reaction:
2NiS(s) + 3O2(g)  2SO2(g) + 2NiO(s)
−149 J/K

Standard Entropy Values
Generally, the more complex the molecule, the higher the
standard entropy value.
 H2O(g) has a higher entropy value than H2 due to the
difference in molecular structure. Because water is
nonlinear, triatomic molecule, it has more rotational and
vibrational motions than the diatomic H2 molecule.

Free Energy and Chemical
Reactions
17.6

Free Energy in a Reaction
Free energy cannot be measured directly for a reaction.
 Free energy must be calculated from other measure
quantities.
 ΔG°=ΔH°−TΔS°
 ΔH° and ΔS° can be solved with known values and
then the equation above can be used to solve for free
energy.

Example Problem
Calculate ΔH°, ΔS° and ΔG° using enthalpy and entropy
values in the Appendix for the following reaction:
2SO2(g) + O2(g)  2SO3(g)
ΔH° = -198 kJ, ΔS° = −187J/K, ΔG° = −142 kJ

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