Chapter 6 Thermochemistry: The Fire Within

1
Chapter 6 Thermochemistry The Fire Within
Study of the heat released or required by
chemical reactions.

• This canopy walkway and net in the Peruvian
  rainforest allow biomass research to be conducted
  high above the ground. The source of most of our
  energy is the Sun. Through the process of
  photosynthesis, solar radiation causes chemical
  reactions in green plants that store the energy
  for future use. We make use of the energy
  captured by plants when we burn fuels. Fossil
  fuels contain energy that has been stored for
  thousands of years, but research into alternative
  fuels is finding ways to make efficient use of
  plants for fuel. This chapter presents the basic
  concepts used in research into the energy changes
  that accompany all chemical reactions.

2
Assignment for Chapter 6
14, 23, 35, 47, 55, 61
3
Energy, Heat, Enthalpy

• CH4(g)2O2(g)?CO2(g)2H2O(l)energy
• C6H12O6(aq)O2(g)?6CO2(g)6H2O(l)energy
• Energetics How energy is transformed and used
• Bioenergetics Study of the use of energy in
  organisms
• Conservation and transformation of energy
• Heat
• Enthalpy

4
Figure 6.1 The potential energy of a mass, m, is
proportional to its height, h, above a surface
that is taken to correspond to zero potential
energy.
5
Figure 6.2 Kinetic energy (represented by the
height of the dark green bar) and potential
energy (the light green bar) can be converted
into one another. However, their sum (the total
height of the bar) is a constant in the absence
of external influences, such as air resistance. A
ball thrown up from the ground loses kinetic
energy as it slows but gains potential energy.
The reverse happens as it falls back to Earth.
Conservation of energy EkEpconst. for an
isolated body
Internal Energy Usum of the kinetic and
potential energies of all the atoms and molecules
in a sample.
6
Figure 6.3In thermodynamics, the world is
divided into a system (the object of interest)
and the surroundings (everything else). In
practice, the surroundings may be a
constant-temperature water bath. The arrows
represent the energy being transferred between
the system and its surroundings.
Heat(energy) exchange(transfer) Matter
exchange(transfer)
7
Figure 6.4 We classify systems into one of three
kinds, according to their interactions with their
surroundings. An open system can exchange matter
and energy with its surroundings. A closed system
can exchange energy but not matter. An isolated
system can exchange neither matter nor energy.
8
Figure 6.5 The booster rockets on the space
shuttle form an open system. The stream of gases
produced by the chemical reaction pours out of
the engines and moves the rocket. (See Applying
Chemistry Case Study 20.)
9
Figure 6.6 When we heat a system, we make use of
a difference in temperature between it and the
surroundings to induce energy to flow through the
walls of the system. Heat flows from high
temperature to low.
Heat is a transfer of energy that occurs as a
result of a temperature difference.
1 cal 4.184 J
10
Figure 6.7 When energy leaves a system as a
result of a temperature difference between the
system and the surroundings, we say that the
system has lost energy as heat. This transfer of
energy stimulates the thermal motion of molecules
in the surroundings.
Thermal motion random molecular motion (in
gases, liquids solids)
11
Figure 6.8 A system does work when it expands
against an external pressure. Here we see a gas
that pushes a piston out against a pressure, P.
We shall see shortly that the work done is
proportional to both the pressure and the change
in volume that the system undergoes.
Work is a transfer of energy that takes place
when an object is moved against an opposing force.
12
Figure 6.9 When a system expands, it performs
work on its surroundings by forcing all the
molecules in another object in the surroundings
to move in the same direction. Here we see the
expansion of a gas raise a weight. The expansion
of gases in the cylinders of automobiles do work
by pushing on the piston, which turns the gears
that move the vehicle.
13
Three Ways for Changing the Internal Energy of a
System

• Adding matter
• Supplying energy as heat (thermal motion)
• Doing work (uniform motion)

14
Figure 6.10 When we wind a spring, the potential
energy of the atoms changes because they are
squashed together and repel one another. The
internal energy of the spring rises as a result
of this increase in potential energy.
Another type of potential energy Interactions
between atoms in a system.
15
Figure 6.11 The internal energy of a system can
be changed either by doing work or by heating.
The diagram shows that the change in internal
energy is positive (U increases) when energy is
supplied in either way. When energy leaves the
system as heat or work, the internal energy falls
and DU is negative.
16
Figure 6.12 The inventor of this elaborate
device, the Keely motor, claimed that it could
generate high pressures and energies from a small
amount of water. However, like other perpetual
motion machines, the Keely motor was found to be
a fraud. The work was accomplished by compressed
air from a hidden source.
17
Figure 6.13 A reaction does work when it
generates a gas. For example, carbon dioxide is
formed from the thermal decomposition of calcium
carbonate. As the gas is formed, it drives back
the surrounding atmosphere.
W?
18
Figure 6.14 The gas in the piston expands a
distance, l, against a pressure, P. The volume
increase is l ? A, where A is the cross-sectional
area of the piston.
19
An Example for Calculating Work

• A gas expands by 12.0 L against a pressure of
  2.0 atm. How much work is done?

20
Figure 6.15 When a reaction (such as the thermal
decomposition of calcium carbonate) takes place
in a closed, constant-volume container, the gas
fills the container but cannot expand against the
surrounding atmosphere. As a result, it does no
work on the surroundings.
21
Figure 6.16 When a system is free to expand
against an external pressure, some of the energy
supplied to it as heat escapes back into the
surroundings as work. As a result, the change in
internal energy is less than the energy supplied.
22
Enthalpy
Change in enthalpy of a system is equal to the
heat transferred to it at constant pressure
23
Figure 6.17 The thermite reaction is another
highly exothermic reactionone that can melt the
metal it produces. In this reaction, aluminum
metal is reacting with iron(III) oxide, Fe2O3,
causing a shower of molten iron sparks. In an
exothermic reaction, energy is lost as heat, the
amount lost depending on the amount of reactants
available.
24
Figure 6.18 The reaction between ammonium
thiocyanate, NH4SCN, and barium hydroxide
octahydrate, Ba(OH)28H2O, absorbs a lot of heat
and can cause water vapor in the air to freeze on
the outside of the beaker. In an endothermic
reaction, energy is absorbed as heat.
2NH4SCN Ba(OH)28H2O?
Ba(SCN)22NH3 (g)10H2O (g)
25
Classroom Exercise

• In an endothermic reaction, 5 kcal of heat is
  absorbed and 1.2 L of gas is generated. How much
  is the change of the internal energy?

26
Figure 6.19 Heat capacity is an extensive
property, so a large object (bottom) has a larger
heat capacity than a small object (top) made of
the same material.
27
Specific Heat Capacity
28
(No Transcript)
29
Using Specific Heat Capacity

• 5.10 g sample of an alloy with specific heat
  capacity of 0.124 J/C/g is heated from 24.2 C to
  138.5 C. How much energy is supplied?

30
Figure 6.20 A pyrotechnics expert attaches fuses
containing potassium chlorate to fireworks set up
in mortars. The highly exothermic reactions of
fireworks do not begin until they are intiated by
a fuse. The fuse ignites a mixture of carbon and
other reducing agents that react with an oxidizer
such as potassium perchlorate.
31
Figure 6.21 The quantity of heat released or
absorbed by a reaction can be measured in this
primitive version of a calorimeter. The outer
polystyrene cup acts as an extra layer of
insulation to ensure that no heat enters or
leaves the inner cup. The quantity of heat
released or absorbed is proportional to the
change in temperature of the calorimeter.
32
Figure 6.22 A bomb calorimeter. The combustion
is initiated with an electrical fuse. Once the
reaction has begun, energy is released as heat
that spreads through the walls of the bomb into
the water. The heat released is proportional to
the temperature change of the entire calorimeter
assembly.
33
Calorimetry

• 50.0 g water at 20 C is mixed with 21 g of iron
  at 90.2 C. The equilibrium temperature is 23.2 C.
  Find the specific heat capacity of iron.

34
Enthalpy H U PV
J. Willard Gibbs
Benoit Paul Émile Clapeyron
Heike Kamerlingh Onnes
Rudolf Clausius
35
Thermodynamics related to eating/drinking, so
much, more than entire physics.
36
How does power come from?
HUPV
37
Ice? cold water ? hot water ? vapor
38
Vaporization enthalpy of vaporization
39
Figure 6.23 Melting (fusion) is an endothermic
process. As molecules acquire energy, they begin
to struggle past their neighbors. Finally the
sample changes from a solid with ordered
molecules (left) to a liquid with disordered,
mobile molecules (right).
Melting/freezing enthalpy of fusion/freezing
40
Figure 6.24 The enthalpy change for the reverse
of a process is the negative of the enthalpy
change for the forward process at the same
temperature.
41
Figure 6.25 The polar ice caps on Mars extend
and recede with the seasons. They are solid
carbon dioxide and form by direct conversion of
the gas to a solid. They disappear by
sublimation. Although some water ice is also
present in the polar caps, the temperature on
Mars never becomes high enough to melt it. On
Mars, ice is just another rock.
42
Figure 6.26 Because enthalpy is a state
property, the enthalpy of sublimation at a given
temperature can be expressed as the sum of the
enthalpies of fusion and vaporization measured at
the same temperature.
Sublimation enthalpy of sublimation
43
Thermochemistry of Some Physical Changes
Vaporization enthalpy of vaporization
Melting/freezing enthalpy of fusion/freezing
Sublimation enthalpy of sublimation
44
(No Transcript)
45
Figure 6.27 The heating curve of water. The
temperature of the solid rises as heat is
supplied. At the melting point, the temperature
remains constant and the heat is used to melt the
sample. When enough heat has been supplied to
melt all the solid, the temperature of the liquid
begins to rise again. A similar pause in the
temperature rise occurs at the boiling point.
46
The Enthalpy of Chemical Change Reaction
Enthalpies

• Reactants?Porductsheat (exothermic)
• Reactantsheat?Products (endothermic)

CH4(g)2O2(g)?CO2(g)2H2O(l)heat CH4(g)2O2(g)?CO
2(g)2H2O(l) ?Hr-890.0 kJ/mol
(Thermochemical Equation)
Molar Reaction Enthalpy (CH4)
47
Figure 6.28 A biophysicist monitors an
experimental fermentation chamber in which fuel
ethanol is being produced from waste biomass by a
genetically engineered strain of bacteria.
48
(No Transcript)
49
(No Transcript)
50
Figure 6.29 This diagram shows how the value of
the reaction enthalpy depends on the physical
states of a product. When water is produced as a
vapor rather than as a liquid in the combustion
of methane, 88 kJ remains stored in the system
for every 2 mol H2O produced.
CH4(g)2O2(g)?CO2(g)2H2O(g) ?Hr-802.0
kJ/mol CH4(g)2O2(g)?CO2(g)2H2O(l) ?Hr-890.0
kJ/mol
51
Figure 6.30 The standard reaction enthalpy is
the difference in enthalpy between the pure
products, each at 1 atm, and the pure reactants
at the same pressure and the specified
temperature (which is commonly but not
necessarily 25C). The scheme here is for the
combustion of methane.
Standard reaction enthalpy CH4(g)2O2(g)?CO2(g)2
H2O(l) ?Ho-890.0 kJ/mol
52
Investigating Matter 6.1 (a)Because fossil fuel
reserves are limited, they must be extracted
wherever they are found. This platform is used to
pump petroleum from beneath the ocean however,
the natural gas accompanying it cannot be easily
transported and so is burned off.
53
Investigating Matter 6.1 (b)An agricultural
researcher assesses the growth rate of a
seedling. Plant photosynthesis is only about 3
efficient, and conditions that increase this
efficiency are actively being sought.
54
Reaction Enthalpy of A Reverse Reaction
P4(s)6Cl2(g)?4PCl3(l)
4PCl3(l) ?P4(s)6Cl2(g)
C6H12O6(aq)6O2(g)?6CO2(g)6H2O(l)
6CO2(g)6H2O(l) ?C6H12O6(aq)6O2(g)
55
Figure 6.31 If the overall reaction can be
broken down into a series of steps, then the
corresponding overall reaction enthalpy is the
sum of the reaction enthalpies of the steps on
the alternative path. None of the steps need be a
reaction that can actually be carried out in the
laboratory.
Hesss Law
Germain Henri Hess (1802 - 1850)
56
Using Hesss Law

• C(s)O2(g)?CO2(g)

C(s)(1/2)O2(g)?CO(g)
CO(g)(1/2)O2(g)?CO2(g)

C(s)O2(g)?CO2(g)
57
Using Hesss Law

• 3C(s)4H2(g)?C3H8(g) (X)

C3H8(g)5O2(g)?3CO2(g)4H2O(l) (A)
C(s)O2(g)?CO2(g) (B)
H2(g)(1/2)O2(g)?H2O(g) (C)
X3B-A4C
58
Figure 6.32 The amount of heat produced or
absorbed in a chemical reaction can be determined
from the reaction stoichiometry.
59
The Enthalpy of Chemical Change Reaction
Enthalpies

• Enthalpy of combustion

Enthalpy of formation
60
CH4(g)2O2(g)?CO2(g)2H2O(g) ?Hc-802.0 kJ/mol
Enthalpy of Combustion
How much heat is produced by burning 150.0 g of
methane?
61
(No Transcript)
62
Figure 6.33 During World War II, fuel was in
short supply and all manner of ingenious
solutions were sought. However, as we can see
from this photograph of a vehicle powered by coal
gas (a mixture of carbon monoxide and hydrogen)
in London, the low enthalpy density of gases
creates storage problems. A modern approach to
using gases to power a vehicle can be seen in
Applying Chemistry Case Study 18.
63
Different Units for Enthalpy
Specific Enthalpy
Enthalpy Density
64
(No Transcript)
65
Figure 6.34 The range and speed of this
electric-powered car depend on the type of
battery it uses. For example, metal-hydride
devices have a longer range than lead-acid
storage batteries. As this driver would agree,
recharging is generally a slow process. The cable
attached to this car may look like a gasoline
hose, but it is actually delivering electricity
while its owner waits. The use of hydrogen as
fuel could reduce the number of refueling stops.
66
Standard Enthalpies of Formation

• The standard enthalpy of formation of an
  element in its most stable form is 0.

67
Figure 6.35 The reaction enthalpy can be
constructed from enthalpies of formation by
imagining the formation of both the reactants and
the products from their respective elements. The
reaction enthalpy is the difference between the
two.
68
Standard Enthalpies of Formation

• 4C(s)6H2(g) O2(g)?2C2H5OH(l) ?Hc-555.38 kJ

69
Classroom Exercise

• 2C(s)O2(g)?2CO (g) ?Hc-221.06 kJ

Standard enthalpies of formation of CO?
70
The standard enthalpy of formation of an
element in its most stable form is 0.
71
Using Standard Enthalpy of Formation
72
Fun Chemistry Calorie of Food

• It is equal to specific combustion enthalpy
  kcal/g
• It is in the sense of average and approximation,
  depending on the location and growth conditions
  of the original produces, the genetics of the
  people who consumes the food

Compound A O2 ? Compound B Compound C
?Hc calorie
It is assumed that the calorie of a compound
(protein, carbonhydrate etc) metabolized in body
is the same as that when the compound is
burnt outside.
73
Measuring Calories

• The caloric value of food is the energy produced
  by combustion of its proteins, carbohydrates and
  fats. The amount of energy liberated by the
  catabolism of food in the body is almost the same
  as the amount liberated when food is burnt
  outside the body. The energy liberated by
  catabolic processes is used for maintaining body
  functions namely digestion, thermoregulation,
  muscular contraction and nerve impulses
  conduction. The amount of energy liberated / unit
  time is the metabolic rate. When food is burnt
  outside the body, all the energy is liberated as
  heat.
• The standard unit of heat energy is the calorie
  (cal), which is defined as the amount of heat
  energy necessary to raise the temperature of 1 ml
  of water by one degree, from 15 to 16 celsius at
  rest.
• A slightly different calorie is used in
  engineering, the international calorie, which
  equals 1/860 international watt-hour (4.1868 J).
  A large calorie, or kilocalorie, usually referred
  to simply as a calorie and sometimes as a
  kilogram calorie, equals 1,000 calories and is
  the unit used to express the energy-producing
  value of food in the calculation of diets.
• The energy released by combustion of foodstuffs
  outside the body can be measured directly and
  indirectly. In direct calorimetry method the
  liberated energy can be measured using a bomb
  calorimeter. It is a metal vessel surrounded by a
  water insulated container. The food is ignited by
  an electric spark. The change in the temperature
  of water is a measure of the calories produced.
• In indirect method, the energy production can
  also be calculated by
• measuring the amount of oxygen consumed
  for combustion of food.
• The amount of oxygen consumed / unit of
  time is proportionate to the
• energy liberated. This method of energy
  estimation is called wet
• combustion. The caloric value of
  carbohydrate is 4.1 Kcal /g,
• protein is 5.65 Kcal/ g and fat is 9.4
  Kcal/g.

74
Case Study 6 (a)Regular exercise not only is
good for the metabolism, it can be fun, too, when
we make it a part of daily life.
When you exercise, you burn the nutrients in your
body by speeding up metabolism and spend more
calories.
75
Case Study 6 (b)Energy consumed (in kilojoules
per hour) in typical activities blue for a 70-kg
male and pink for a 58-kg female.
76
(No Transcript)
77
Biological systems are wonderful chemicals and
wonderful chemical plants
78
Thermodynamics related to eating/drinking, so
much, more than entire physics.
Now, a six-table banquet is served. When enjoying
the delicacies, refresh yourself of the
meaning of calorie Dr. Ding told you.
79
Assignment for Chapter 6
14, 23, 35, 47, 55, 61