Chapter 6

Thermochemistry
 Heating Your Home
• most homes burn fossil fuels to generate heat
• the amount the temperature of your home
increases depends on several factors
✓how much fuel is burned
✓the volume of the house
✓the amount of heat loss
✓the efficiency of the burning process
✓can you think of any others?
Tro, Chemistry: A Molecular Approach 2
 Nature of Energy
• even though Chemistry is the study of
matter, energy effects matter
• energy is anything that has the capacity to
do work
• work is a force acting over a distance
✓Energy = Work = Force x Distance
• energy can be exchanged between objects
through contact
✓collisions
Tro, Chemistry: A Molecular Approach 3
 Classification of
Energy
• Kinetic energy is
energy of motion or
energy that is being
transferred
✓thermal energy is
kinetic

Tro, Chemistry: A Molecular Approach 4

 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

Tro, Chemistry: A Molecular Approach 5

 Law of Conservation of Energy
• energy cannot be created or
destroyed
✓ First Law of
Thermodynamics
• energy can be transferred
between objects
• energy can be transformed
from one form to another
✓ heat → light → sound
Tro, Chemistry: A Molecular Approach 6
 Some Forms of Energy
• Electrical
✓ kinetic energy associated with the flow of electrical charge
• Heat or Thermal Energy
✓ kinetic energy associated with molecular motion
• Light or Radiant Energy
✓ kinetic energy associated with energy transitions in an atom
• Nuclear
✓ potential energy in the nucleus of atoms
• Chemical
✓ potential energy in the attachment of atoms or because of
their position
Tro, Chemistry: A Molecular Approach 7
 Units of Energy
• the amount of kinetic energy an
object has is directly proportional
to its mass and velocity
✓ KE = ½mv2
• when the mass is in kg and
speed in m/s, the unit for kinetic
kg
energy is 2 • m 2

s
• 1 joule of energy is the amount of
energy needed to move a 1 kg mass
at a speed of 1 m/s
kg • m 2
✓ 1J=1 2 8
s
 Units of Energy
• joule (J) is the amount of energy needed to move
a 1 kg mass a distance of 1 meter
✓1 J = 1 N∙m = 1 kg∙m2/s2
• calorie (cal) is the amount of energy needed to
raise one gram of water by 1°C
✓kcal = energy needed to raise 1000 g of water 1°C
✓food Calories = kcals
Energy Conversion Factors
1 calorie (cal) = 4.184 joules (J) (exact)
1 Calorie (Cal) = 1000 calories (cal)
1 kilowatt-hour (kWh) = 3.60 x 106 joules (J)
Tro, Chemistry: A Molecular Approach 9
 Energy Use
Energy Energy Energy
Required to Energy used to Used by
Raise Required to Run 1 Average
Unit
Temperature Light 100-W Mile U.S.
of 1 g of Bulb for 1 hr Citizen in
Water by 1°C (approx) 1 Day
joule (J) 4.18 3.60 x 105 4.2 x 105 9.0 x 108

calorie (cal) 1.00 8.60 x 104 1.0 x 105 2.2 x 108

Calorie (Cal) 0.00100 86.0 100. 2.2 x 105

kWh 1.16 x 10-6 0.100 0.12 2.5 x 102

Tro, Chemistry: A Molecular Approach 10
 Energy Flow and
Conservation of Energy
• we define the system as the material or process we are
studying the energy changes within
• we define the surroundings as everything else in the
universe
• Conservation of Energy requires that the total energy
change in the system and the surrounding must be zero
✓ DEnergyuniverse = 0 = DEnergysystem + DEnergysurroundings
✓ D is the symbol that is used to mean change
➢ final amount – initial amount

11
 Internal Energy
• the internal energy is the total amount of
kinetic and potential energy a system possesses
• 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
✓DE = Efinal – Einitial
✓DEreaction = Eproducts - Ereactants
Tro, Chemistry: A Molecular Approach 12
 State Function

Tro, Chemistry: A Molecular Approach 13

 Energy Diagrams
• energy diagrams are a

Internal Energy
“graphical” way of showing final
the direction of energy flow energy added
during a process DE = +
• if the final condition has a initial
larger amount of internal
energy than the initial
condition, the change in the

Internal Energy
internal energy will be + initial
• if the final condition has a energy removed
smaller amount of internal DE = ─
final
energy than the initial
condition, the change in the
internal energy will be ─
Tro, Chemistry: A Molecular Approach 14
 Energy Flow
• when energy flows out of a
system, it must all flow into Surroundings
the surroundings DE +
• when energy flows out of a System
system, DEsystem is ─ DE ─
• when energy flows into the
surroundings, DEsurroundings is +
• therefore:
─ DEsystem= DEsurroundings

Tro, Chemistry: A Molecular Approach 15

 Energy Flow
• when energy flows into a
system, it must all come from Surroundings
the surroundings DE ─
• when energy flows into a
system, DEsystem is + System
DE +
• when energy flows out of the
surroundings, DEsurroundings is ─
• therefore:
DEsystem= ─ DEsurroundings

Tro, Chemistry: A Molecular Approach 16

 How Is Energy Exchanged?
• 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
DE = q + w
system gains heat energy system releases heat energy
q (heat) + ─
system releases energy by
system gains energy from work
w (work) +
doing work
─
system gains energy system releases energy
DE + ─
Tro, Chemistry: A Molecular Approach 17
 Calculate the change in the internal energy for
a process in which a system absorbs 140 J of
heat from the surroundings and does 85 J of
work on the surroundings.

Tro, Chemistry: A Molecular Approach 18

 Nature of Energy
• even though Chemistry is the study of
matter, energy effects matter
• energy is anything that has the capacity to
do work
• work is a force acting over a distance
✓Energy = Work = Force x Distance
• energy can be exchanged between objects
through contact
✓collisions
Tro, Chemistry: A Molecular Approach 19
 Law of Conservation of Energy
• energy cannot be created or
destroyed
✓ First Law of
Thermodynamics
• energy can be transferred
between objects
• energy can be transformed
from one form to another
✓ heat → light → sound
Tro, Chemistry: A Molecular Approach 20
 First Law of Thermodynamics
• System - the material or process we are
studying the energy changes within
• Surroundings - everything else in the
universe
• An open system - one in which matter
and energy can be exchanged with the
surroundings.
• Closed systems - systems that can
exchange energy but not matter with
their surroundings
• An isolated system - one in which
neither energy nor matter can be
exchanged with the surroundings.

21
 Internal Energy
• the internal energy is the total amount of
kinetic and potential energy a system possesses
• 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
✓DE = Efinal – Einitial
✓DEreaction = Eproducts - Ereactants
Tro, Chemistry: A Molecular Approach 22
 Internal Energy
Sketch the energy diagram representing the
reaction MgCl2 (s) → Mg(s) + Cl2 (g) knowing that
the internal energy for a mixture of Mg(s) and
Cl2(g) is larger than that of MgCl2(s).

Tro, Chemistry: A Molecular Approach 23

 Energy Exchange

• energy is exchanged between the system and

surroundings through either heat exchange or
work being done

Tro, Chemistry: A Molecular Approach 24

Practice Exercise
Relating Heat and Work to Changes of Internal Energy

Gases A(g) and B(g) are confined in a cylinder-and-piston

arrangement like that in the figure given below and react to form
a solid product C(s) :
A(g) + B(g) → C(s)

As the reaction occurs, the system loses 1150 J of heat to the

surroundings. The piston moves downward as the gases react to
form a solid. As the volume of the gas decreases under the
constant pressure of the atmosphere, the surroundings do 480 J
of work on the system. What is the change in the internal energy
of the system?

25
Endothermic and Exothermic Reactions
• when DH is ─, heat is being released by the system
• reactions that release heat are called exothermic reactions
• when DH is +, heat is being absorbed by the system
• reactions that release heat are called endothermic reactions
• chemical heat packs contain iron filings that are oxidized in
an exothermic reaction ─ your hands get warm because the
released heat of the reaction is absorbed by your hands
• chemical cold packs contain NH4NO3 that dissolves in
water in an endothermic process ─ your hands get cold
because they are giving away your heat to the reaction 26
27
28
29
 State Function

Tro, Chemistry: A Molecular Approach 30

State Function
• Internal energy is a state function, but
heat and work are not.
(a) A battery shorted out by a wire
loses energy to the surroundings only
as heat; no work is performed.
(b) A battery discharged through a
motor loses energy as work (to make
the fan turn) and also loses some
energy as heat.

• The value of ΔE is the same for both

processes even though the values of q
and win (a) are different from those in
(b).
31
Give it Some Thought
You lose 5 pounds over a 30-day period. Which of the
following quantities act like a state function: the amount
of calories you consume, your weight, or the number of
calories burned through exercise?

32
 Enthalpy
• the enthalpy, H, of a system is the sum of the internal
energy of the system and the product of pressure and
volume
✓ H is a state function
H = E + PV
• the enthalpy change, DH, of a reaction is the heat
evolved in a reaction at constant pressure
DHreaction = qreaction 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
Tro, Chemistry: A Molecular Approach 33
 Pressure -Volume Work
• PV work is work that is the result of a volume change
against an external pressure
• when gases expand, DV is +, but the system is doing work
on the surroundings so w is ─
• as long as the external pressure is kept constant
─Work = External Pressure x Change in Volume
w = ─PDV
✓ to convert the units to joules use 101.3 J = 1 atm∙L

Tro, Chemistry: A Molecular Approach 34

 Pressure -Volume Work
• PV work is work that is the result of a volume change
against an external pressure
• when gases expand, DV is +, but the system is doing work
on the surroundings so w is ─
• as long as the external pressure is kept constant
─Work = External Pressure x Change in Volume
w = ─PDV
✓ to convert the units to joules use 101.3 J = 1 atm∙L

Tro, Chemistry: A Molecular Approach 35

Give it Some Thought
If a system does not change its volume during the course
of a process, does it do pressure–volume work?

36
Sample Exercise
If a balloon is inflated from 0.100 L to 1.85 L against an
external pressure of 1.00 atm, how much work is done?
Given: V1=0.100 L, V2=1.85 L, P=1.00 atm
Find: w (in J)
Concept Plan:
P, DV w

w = - P • DV
Relationships: 101.3 J = 1 atm L

Solution:
DV = V2 − V1 w = −P • DV 101.3 J
− 1.75 atm • L 
1 atm • L
DV = 1.85 L - 0.100 L = −(1.00 atm ) • (1.75 L )
= −1.75 atm • L = - 177 J
= 1.75 L
Check:
the unit and sign are correct
Practice Exercise
Calculating Pressure-Volume Work

A fuel is burned in a cylinder equipped with a piston. The initial

volume of the cylinder is 0.250 L, and the final volume is
0.980 L.

If the piston expands against a constant pressure of 1.35 atm, how

much work (in J) is done? (1 L-atm = 101.3 J)

38
Enthalpy Change
• the enthalpy change, DH, of a
reaction is the heat evolved in a
reaction at constant pressure
DHreaction = qreaction 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

Tro, Chemistry: A Molecular Approach 39

Practice Exercise
Determining the Sign of DH

Indicate the sign of the enthalpy change, ΔH, in the following

processes carried out under atmospheric pressure and indicate
whether each process is endothermic or exothermic:

(a) An ice cube melts;

(b) 1 g of butane (C4H10) is combusted in sufficient
oxygen to give complete combustion to CO2 and H2O.

40
Practice Exercise
Determining the Sign of DH

Molten gold poured into a mold solidifies at atmospheric

pressure. With the gold defined as the system, is the solidification
an exothermic or endothermic process?

41
 Enthalpies of Reaction
• the enthalpy change for a chemical reaction is given by
ΔH = Hproducts – Hreactants
• The enthalpy change that accompanies a reaction is
called either the enthalpy of reaction or the heat of
reaction.
✓ Sometimes written ΔHrxn

• Balanced chemical equations that show the associated

enthalpy change are called thermochemical equations.

Tro, Chemistry: A Molecular Approach 42

 Enthalpies of Reaction

• Balanced chemical equations that show the associated

enthalpy change are called thermochemical equations.

Tro, Chemistry: A Molecular Approach 43

 Enthalpies of Reaction
The following guidelines are helpful when using
thermochemical equations and enthalpy diagrams:

1. Enthalpy is an extensive property. The magnitude of ΔH is

proportional to the amount of reactant consumed in the process.

Example:
890 kJ of heat is produced when 1 mol of CH4 is burned in a
constant-pressure system:

Tro, Chemistry: A Molecular Approach 44

 Enthalpies of Reaction
The following guidelines are helpful when using
thermochemical equations and enthalpy diagrams:

2. The enthalpy change for a reaction is equal in magnitude,

but opposite in sign, to ΔH for the reverse reaction.

Example:
ΔH for the reverse of Equation:

is +890 kJ:

Tro, Chemistry: A Molecular Approach 45

 Enthalpies of Reaction
The following guidelines are helpful when using
thermochemical equations and enthalpy diagrams:

3. The enthalpy change for a reaction depends on the states of

the reactants and products.

Example:

▪ Original ΔH with liquid water: +890 kJ

▪ Subtract the energy needed to vaporize the water:
+890 kJ − 88 kJ = +802 kJ

Tro, Chemistry: A Molecular Approach 46

Sample Exercise
Relating DH to Quantities of Reactants and Products

How much heat is released when 4.50 g of methane gas is burned

in a constant-pressure system?

Solution:
• Conversion

• o

47
Practice Exercise
Relating DH to Quantities of Reactants and Products

The complete combustion of ethanol, C2H5OH (molar mass=46.0

g/mol), proceeds as follows:

What is the enthalpy change for combustion of 15.0 g of

ethanol?

a) -12.1 kJ
b) -181 kJ
c) -422 kJ
d) -555 kJ
e) -1700 kJ
48
 Calorimetry
• Calorimetry – the measurement of heat flow
• Calorimeter – the device used to measure heat flow

Heat Capacity and Specific Heat

• when a system absorbs heat, its temperature increases
• the increase in temperature is directly proportional to the
amount of heat absorbed
• the proportionality constant is called the heat capacity, C
• The heat capacity of an object is the amount of heat
required to raise its temperature by 1K or 1°C.
• Units are J/(g-°C) or J/g-K

Tro, Chemistry: A Molecular Approach 49

 Heat Capacity and Specific Heat
• The heat capacity of one mole of a substance is
called its molar heat capacity, Cm.
• The heat capacity of one gram of a substance is
called its specific heat capacity, or merely its
specific heat, Cs

Tro, Chemistry: A Molecular Approach 50

 Heat Capacity and Specific Heat
Example:
• Water’s Specific Heat
For example, 209 J is required to
increase the temperature of 50.0 g of
water by 1.00 K.
Thus, the specific heat of water is:

Tro, Chemistry: A Molecular Approach 51

 Heat Capacity and Specific Heat

• Higher Cs → substance heats up slowly

• Lower Cs → substance heats up quickly

• Water’s high specific heat helps stabilize temperatures

in coastal areas.
Tro, Chemistry: A Molecular Approach 52
Sample Exercise
Relating Heat, Temperature Change, and Heat Capacity

(a) How much heat is needed to warm 250 g of water (about 1

cup) from 22 °C (about room temperature) to 98 °C (near its
boiling point)?
(b) What is the molar heat capacity of water?

Solution:
(a) The water undergoes a temperature change of:

53
Sample Exercise
Relating Heat, Temperature Change, and Heat Capacity

(a) How much heat is needed to warm 250 g of water (about 1

cup) from 22 °C (about room temperature) to 98 °C (near its
boiling point)?
(b) What is the molar heat capacity of water?

Solution:
(b) The molar heat capacity is the heat capacity of one mole of
substance. Using the atomic weights of hydrogen and oxygen,
we have: 1 mol H2O = 18.0 g H2O

From the specific heat used in part (a), we have:

54
Practice Exercise
Relating Heat, Temperature Change, and Heat Capacity

(a) Large beds of rocks are used in some solar-heated homes to

store heat. Assume that the specific heat of the rocks is 0.82 J/g-K.
Calculate the quantity of heat absorbed by 50.0 kg of rocks if their
temperature increases by 12.0 °C.

(b) What temperature change would these rocks undergo if they

emitted 450 kJ of heat?

55
Practice Exercise
Relating Heat, Temperature Change, and Heat Capacity

The specific heat of octane, C8H18 (l), is 2.22 J/g-K.

(a) How many J of heat are needed to raise the temperature of

80.0 g of octane from 10.0 to 25.0 °C?

(b) Which will require more heat, increasing the temperature of 1

mol of C8H18 (l) by a certain amount or increasing the temperature
of 1 mol of H2O (l) by the same amount?

56
 Constant-Pressure Calorimetry
• Constant-pressure
calorimetry is a method
used to measure the heat
change (ΔH) of reactions
that occur in solution,
under the constant pressure
of the atmosphere.

• It’s especially useful for

reactions like neutralization
or dissolution.

Tro, Chemistry: A Molecular Approach 57

 Constant-Pressure Calorimetry
• In an exothermic reaction, heat is released by the
system. The water absorbs it, and the temperature
rises.
• In an endothermic reaction, heat is absorbed by the
system. The water loses heat, and the temperature
drops.

• The heat gained or lost by the solution is equal in

magnitude—but opposite in sign—to the heat of the
reaction.

Tro, Chemistry: A Molecular Approach 58

Sample Exercise
Measuring ΔH Using a Coffee-Cup Calorimeter

When a student mixes 50 mL of 1.0 M HCl and 50 mL of 1.0 M

NaOH in a coffee-cup calorimeter, the temperature of the
resultant solution increases from 21.0 to 27.5 °C.

Calculate the enthalpy change for the reaction in kJ/mol HCl,

assuming that the calorimeter loses only a negligible quantity of
heat, that the total volume of the solution is 100 mL, that its
density is 1.0 g/mL, and that its specific heat is 4.18 J/g-K.

59
Sample Exercise
Measuring ΔH Using a Coffee-Cup Calorimeter

Solution:
Because the total volume of the solution is 100 mL, its mass is:
(100 mL)(1.0 g/mL) = 100 g

The temperature change is:

60
Practice Exercise
Measuring ΔH Using a Coffee-Cup Calorimeter

When 50.0 mL of 0.100 M AgNO3 and 50.0 mL of 0.100 M HCl

are mixed in a constant-pressure calorimeter, the temperature
of the mixture increases from 22.30 to 23.11 °C. The temperature
increase is caused by the following reaction:

61
 Constant-Volume Calorimetry
• Constant-volume
calorimetry is a technique
used to measure the heat
released during a reaction
that occurs in a sealed,
rigid container—called a
bomb calorimeter.

• measuring the change in

internal energy, or ΔE,
rather than enthalpy (ΔH)

Tro, Chemistry: A Molecular Approach 62

 Measuring DE,
Calorimetry at Constant Volume
• since DE = q + w, we can determine DE by measuring q and w
• in practice, it is easiest to do a process in such a way that there is
no change in volume, w = 0
✓ at constant volume, DEsystem = qsystem
• in practice, it is not possible to observe the temperature changes
of the individual chemicals involved in a reaction – so instead,
we use an insulated, controlled surroundings and measure the
temperature change in it
• the surroundings is called a bomb calorimeter and is usually
made of a sealed, insulated container filled with water
qsurroundings = qcalorimeter = ─qsystem
─DEreaction = qcal = Ccal x DT
Tro, Chemistry: A Molecular Approach 63
Sample Exercise
Measuring qrxn Using a Bomb Calorimeter

The combustion of methylhydrazine (CH6N2), a liquid rocket fuel,

produces N2 (g), CO2 (g), and H2O(l) :

When 4.00 g of methylhydrazine is combusted in a bomb

calorimeter, the temperature of the calorimeter increases from
25.00 to 39.50 °C. In a separate experiment the heat capacity of
the calorimeter is measured to be 7.794 kJ/°C.

Calculate the heat of reaction for the combustion of a mole of

CH6N2.

64
Sample Exercise
Measuring qrxn Using a Bomb Calorimeter

Solution:
• For combustion of the 4.00-g sample of methylhydrazine, the
temperature change of the calorimeter is:

• We can use ΔT and the value for Ccal to calculate the heat of
reaction

• We can readily convert this value to the heat of reaction for a

mole of CH6N2:

65
Practice Exercise
Measuring qrxn Using a Bomb Calorimeter

66
The Regulation of Body Temperature
• The foods, such as glucose,
are metabolized—a process
that is essentially controlled
oxidation to CO2 and H2O:

• Perspiration is predominantly water, so the process is the

endothermic conversion of liquid water into water vapor:
 Relationships Involving DHrxn
• when reaction is multiplied by a factor, DHrxn is
multiplied by that factor
✓because DHrxn is extensive
C(s) + O2(g) → CO2(g) DH = -393.5 kJ
2 C(s) + 2 O2(g) → 2 CO2(g) DH = 2(-393.5 kJ) = -787.0 kJ
• if a reaction is reversed, then the sign of DH is
reversed
CO2(g) → C(s) + O2(g) DH = +393.5 kJ

Tro, Chemistry: A Molecular Approach 68

 Relationships Involving DHrxn
Hess’s Law
• if a reaction can be
expressed as a series
of steps, then the
DHrxn for the overall
reaction is the sum of
the heats of reaction
for each step

Tro, Chemistry: A Molecular Approach 69

 Sample – Hess’s Law
Given the following information:
2 NO(g) + O2(g) → 2 NO2(g) DH° = -173 kJ
2 N2(g) + 5 O2(g) + 2 H2O(l) → 4 HNO3(aq) DH° = -255 kJ
N2(g) + O2(g) → 2 NO(g) DH° = +181 kJ
Calculate the DH° for the reaction below:
3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g) DH° = ?

[32 NO2(g) → 32 NO(g) + 1.5 O2(g)]

O2(g)] x 1.5 DH° = (+259.5
1.5(+173kJ)
kJ)
[1
[2 N2(g) + 2.5
5 OO 2(g)
2(g) + +2 1HH →→4 2HNO
2O(l)
2O(l) HNO x 0.5 DH° = (-128
3(aq)]
3(aq)] kJ) kJ)
0.5(-255
[2 NO(g) → N2(g) + O2(g)] DH° = -181 kJ
3 NO2(g) + H2O(l) → 2 HNO3(aq) + NO(g) DH° = - 49 kJ

Tro, Chemistry: A Molecular Approach 70

Practice Exercise
Hess’s Law

71
 Enthalpies of Formation
• Enthalpy of Formation, ΔHf
- the heat change when 1 mole of a compound is
formed from its elements with all substances
• Standard Enthalpy of Formation, ΔHf
- heat change when 1 mole of a compound is
formed from its elements with all substances in
their standard states.
- Standard State: the most stable physical form of
a substance at 1 atmosphere of pressure and a
specified temperature—usually 25°C (298 K)
Tro, Chemistry: A Molecular Approach 72
 Enthalpies of Formation

Tro, Chemistry: A Molecular Approach 73

 Standard Enthalpy of Formation

• If an element exists in more than one form

under standard conditions, the most stable form
of the element is usually used for the formation
reaction.

Tro, Chemistry: A Molecular Approach 74

 Standard Enthalpy of Formation

Tro, Chemistry: A Molecular Approach 75

 Standard Enthalpy of Formation
Writing heat of formation reactions:
• Balance the equation so that one mole of the
compound is produced.
• Remember the diatomic molecules and write them
correctly (H2, N2, O2, F2, Cl2, Br2, I2).
• The reactants must be elements, not polyatomic
ions.

Tro, Chemistry: A Molecular Approach 76

Sample Exercise
Equations Associated with Enthalpies of Formation

77
Practice Exercise
Equations Associated with Enthalpies of Formation

78
 Enthalpies of Formation
Using Enthalpies of Formation to Calculate Enthalpies of Reaction

We can write this equation as the sum of three equations

associated with standard enthalpies of formation:
 Enthalpy Diagram for Propane Combustion

Tro, Chemistry: A Molecular Approach 80

 Enthalpies of Formation
Using Enthalpies of Formation to Calculate Enthalpies of Reaction

The standard enthalpy change of a reaction is the sum of the

standard enthalpies of formation of the products minus the
standard enthalpies of formation of the reactants:
Sample Exercise
Calculating an Enthalpy of Reaction from Enthalpies of Formation

Solution:

82
Practice Exercise
Calculating an Enthalpy of Reaction from Enthalpies of Formation

Use the provided table to calculate the enthalpy change for the
combustion of 1 mol of ethanol:

83
Practice Exercise
Calculating an Enthalpy of Reaction from Enthalpies of Formation

Use the provided table to calculate the enthalpy change for the
combustion of 1 mol of glucose:

84
Practice Exercise
Calculating an Enthalpy of Reaction from Enthalpies of Formation

Calculate the standard heat of reaction for the reaction of

nitrogen monoxide gas with oxygen to form nitrogen dioxide gas.

85
 Bond Enthalpies
Bond enthalpy - refers to the amount of energy stored in the
chemical bonds between any two atoms in a molecule.

The net change in energy during a chemical reaction is the

difference between:
• how much energy it takes to break chemical bonds; and
• how much energy is released when bonds form.

The energy change will equal:

Amount of energy required Amount of energy released

to break the bonds of ¯ when the bonds of the
the reactant molecules products form
 Bond Enthalpies
The dissociation of Cl2(g) into chlorine atoms results when the
Cl—Cl bond is broken:

• The bond enthalpy for a Cl—Cl bond is equal to the enthalpy

of this reaction, 242 kJ/mol.
• We use the letter D followed by the bond in question to
represent bond enthalpies.
✓ D(Cl—Cl) is the bond enthalpy for the Cl2 bond; and
✓ D(H—Br) is the bond enthalpy for the HBr bond
 Bond Enthalpies
Many important bonds, such as the C—H bond, exist only in
polyatomic molecules.

For these bonds, we usually use average bond enthalpies.

Bond Enthalpies
Bond Enthalpies
 Bond Enthalpies
The oxygen atoms in an O2 molecule are held together by a
double bond, O=O, rather than a single bond, O—O.
 Bond Enthalpies
Bond Enthalpies and the Enthalpies of Reactions

• Breaking bonds is always endothermic; and

• Forming bonds is always exothermic.

Strategies for estimating reaction enthalpies:

1. Supply enough energy to break those bonds in the

reactants that are not present in the products.

2. Form the bonds in the products that were not present in

the reactants.
 Bond Enthalpies
Bond Enthalpies and the Enthalpies of Reactions

The enthalpy of the reaction, ΔHrxn , is estimated as the sum of

the bond enthalpies of the bonds broken minus the sum of the
bond enthalpies of the bonds formed:

•If the total enthalpy of the broken bonds is larger, the reaction
is endothermic (ΔHrxn > 0);
•if the total enthalpy of the newly formed bonds is larger, the
reaction is exothermic (ΔHrxn < 0)
 Bond Enthalpies
Bond Enthalpies and the Enthalpies of Reactions

The gas-phase reaction between methane, CH4, and chlorine

to produce methyl chloride, CH3Cl, and hydrogen chloride, HCl:

Bonds broken: 1 mol C—H, 1 mol Cl—Cl

Bonds formed: 1 mole C—Cl, 1 mol H—Cl
 Bond Enthalpies
Bond Enthalpies and the Enthalpies of Reactions

The gas-phase reaction between methane, CH4, and chlorine

to produce methyl chloride, CH3Cl, and hydrogen chloride, HCl:

Bonds broken: 1 mol C—H, 1 mol Cl—Cl

Bonds formed: 1 mole C—Cl, 1 mol H—Cl
Sample Exercise
Estimating Reaction Enthalpies from Bond Enthalpies

Estimate ΔH for the following combustion reaction.

Solution:

96
Practice Exercise
Estimating Reaction Enthalpies from Bond Enthalpies

Use the average bond enthalpies in the table

provided below to estimate ΔH for the
combustion of ethanol.

97
 Foods and Fuel
Fuel Value - energy released when one gram of any substance is
combusted; can be measured by calorimetry

Foods
• Most of the energy our bodies need comes from
carbohydrates and fats.
• Glucose is transported by the blood to cells where it reacts
with O2 in a series of steps, eventually producing CO2 (g),
H2O(l), and energy:
 Foods and Fuel
Foods
• The reaction of tristearin, C57H110O6, a typical fat

Fats are well suited to serve as the body’s energy reserve for
at least two reasons:
1. They are insoluble in water
2. they produce more energy per gram than either proteins
or carbohydrates

The average fuel value of fats is 38 kJ/g (9 kcal/g).

 Foods and Fuel
Foods
• The combustion of carbohydrates and fats in a bomb
calorimeter gives the same products as when they are
metabolized in the body.
• Proteins contain nitrogen, which is released in the bomb
calorimeter as N2. In the body this nitrogen ends up
mainly as urea, (NH2)2CO.
• Proteins are used by the body mainly as building materials
for organ walls, skin, hair, muscle, and so forth.
• On average, the metabolism of proteins produces 17 kJ/g
(4 kcal/g), the same as for carbohydrates.
Tro, Chemistry: A Molecular Approach 101
Sample Exercise 1
Estimating the Fuel Value of a Food from Its Composition

a) A 28-g (1-oz) serving of a popular breakfast cereal served with

120 mL of skim milk provides 8 g protein, 26 g carbohydrates,
and 2 g fat. Using the average fuel values of these substances,
estimate the fuel value (caloric content) of this serving.
Solution:

102
Sample Exercise 2
Estimating the Fuel Value of a Food from Its Composition

b) A person of average weight uses about 100 Cal/mi when

running or jogging. How many servings of this cereal provide
the fuel value requirements to run 3 mi?

Solution:

103
 Foods and Fuel
Fuel
When fuels burn completely:
•Carbon → CO₂
•Hydrogen → H₂O

Fuels with more carbon and hydrogen (like coal or propane)

have higher fuel values than those with less (like wood).
 Foods and Fuel
Fuel
When fuels burn completely:
•Carbon → CO₂
•Hydrogen → H₂O

Fuels with more carbon and hydrogen (like coal or propane)

have higher fuel values than those with less (like wood).
 Foods and Fuel
Other Energy Sources

• Nonrenewable Energy
Nuclear energy - the energy released in either the fission
(splitting) or the fusion (combining) of atomic nuclei; free
of the polluting emissions that are a major problem with
fossil fuels.

• Renewable energy sources include:

➢ Solar energy from the Sun,
➢ Wind energy harnessed by windmills
➢ Geothermal energy from the heat stored inside Earth,
➢ Hydroelectric energy from flowing rivers, and biomass
energy from crops and biological waste matter.