Chapter 7: Energy and Chemical
Change
• Energy is the ability to do work and supply
heat
• Work is motion against an opposing force
Kinetic Energy  KE  12 mv2
• Potential energy (PE) is the energy of
position or internal arrangement
• KE can be converted into PE and vice versa
When the child is at points (a) and
(c) they have only PE; at point (b)
only KE. Total energy is conserved
(PE + KE = constant).
When fully
compressed or
extended only
PE; at natural
length only KE.
Total energy is
conserved.
• The SI unit of energy is the joule (J)
• A 2 kg object moving at 1 meter per second
has 1 J of kinetic energy
• You may also encounter the calorie (cal)
1 J  (2 kg) 
1
2

1m 2
1s
 1 kg m 2 s -2
1 cal  4.184 J (exactly)
• The energy that is transferred as heat comes
from the object’s internal energy
• The energy associated with the motion of
the object’s molecules is referred to as its
molecular kinetic energy
• The internal energy is often given the
symbol E or U
• We are interested in the change in E:
E  Efinal - Einitial or E  Eproducts - Ereactants
• The temperature of an object is related to
the average kinetic energy of its atoms and
molecules
The temperature for curve
(1) is lower than for curve
(2) because the average
kinetic energy is lower.
•
Heat is a transfer of energy due to a
temperature difference
• The object we are interested in is called the
system
• Everything outside the system is called the
surroundings
• A boundary separates the system from the
surroundings
– Open systems can gain or lose mass and
energy across their boundaries
– Closed systems can absorb or release energy,
but not mass, across their boundaries
– Isolated systems cannot exchange energy or
matter with their surroundings
– When heat is gained by an object, it is written
as a positive number
– When heat is lost by an object, it is written as a
negative number
• A spontaneous change is one that continues
on its own
– Heat flows spontaneously from a warmer to
colder object
• The heat directly gained or lost by an object
is directly proportional to the temperature
change it undergoes
• The object’s specific heat (C) relates the
heat (q) to the objects temperature change
q  C (tfinal - tinitial )  C t
• The heat capacity is the amount of heat
needed to raise the object’s temperature by
one degree Celsius and has the units J/°C
• C is an extensive property that can be
determined from experiment, and is
proportional to the sample mass
• The specific heat capacity (s) is an
intensive property, and is unique for each
substance
C  m s or s  C/m with units of J g -1 C-1
• For example
Specific Heat
Substance J g -1 C -1 (25 C)
Copper
0.387
Gold
Silver
0.129
0.235
Water
4.18
A large specific heat means
the substance releases a large
amount of heat as it cools.
The heat absorbed or released
by an substance is then:
q  m s t
Example: The temperature of 251 g of water is
changed from 25.0 to 30.0 °C. How much heat
was transferred to the water?
ANALYSIS: Connect heat to the temperature
change.
SOLUTION:
q  m s t
 (251 g)  (4.18 J g -1 C-1 )  (30.0 - 25.0)C
 5250 J  5.25 kJ
Note: Heat was absorbed because q is positive
• Chemical bonds are the net attractive force
between nuclei and electrons in compounds
• Breaking a chemical bond requires energy
The attraction of the electrons for
the nuclei in the hydrogen
molecule is strong enough to
overcome the nucleus-nucleus and
electron-electron repulsions.
• Making a chemical bond releases energy
• The potential energy that resides in
chemical bonds is called chemical energy
• Chemical reactions generally involve both
making and breaking chemical bonds
• The net gain or loss of energy is often in the
form of heat
• Any reaction where heat is a product is
called exothermic
• Reactions that consume energy are called
endothermic
• Reactions can release heat by replacing
“weak” bonds with “strong” ones
• The amount of heat absorbed or released by
a chemical reaction is called the heat of
reaction
• A calorimeter can be used to measure the
heat of reaction
• Calorimeters are usually designed to
measure heats of reaction under conditions
of constant volume or constant pressure
• Pressure is the amount of force acting on a
unit area: pressure  force
area
• Atmospheric pressure is the pressure
exerted by the mixture of gases in the
atmosphere
• At sea level the atmospheric pressure is
about 14.7 lb/in²
• Other common pressure units are the
atmosphere (atm) and bar:
14.696 lb/in² = 1.0000 atm = 1.0133 bar
• qv and qp are used to show heats measured at
constant volume or pressure, respectfully
• In reactions where gases are produce or
consumed qv and qp can be very different
Pressure-volume work
(a) A gas confined under
pressure. (b) The gas does
pressure-volume work on
the surroundings when it
expands.
• If the volume change is V , work (w) is
w   PV where P is the opposing pressure
• Note that the work of expansion is negative
• Work and heat are alternate ways to transfer
energy
• Their sum is the change in internal energy
the system undergoes
E  q  w
• This is a statement of the first law of
thermodynamics, which says that energy
cannot be created or destroyed
• Heat and work are not state functions
because they depend on the path between
the final and initial state
Heat and work depend
on the path. Both paths
give the same value for
the internal energy
change. In path 1, this
appears as heat (q). In
path 2 this appears
mostly as work (w).
• The heat produced by a combustion reaction
is called the heat of combustion
• The heats are measured in closed containers
because the reactions consume and produce
gases
• The instrument used to measure these heats
is called a bomb calorimeter
• The reaction is run at constant volume so
that
E  qv
A bomb calorimeter. The
reaction chamber has a constant
volume so no work is done. The
heat released by the combustion
is absorbed by the “bomb” and
surrounding water.
• Heats of reactions in solution are usually
run in open containers at constant pressure
• They may transfer heat and expansion work
• The heat change measured at constant
pressure is the enthalpy, H
H products-H reactants  H  E PV  qp
• Enthalpy is also a state function
– H is negative for an exothermic process
– H is positive for an endothermic process
• The the difference in the values of the
internal energy and enthalpy change can be
large for reactions that consume or release
gases
A coffee cup calorimeter can be used to
measure heats of reaction at constant
pressure. Heat can be released or absorbed,
resulting in a change in temperature of the
solution.
• The amount of heat that a reaction produces
or absorbs depends on the number of moles
of reactant that react
• A set of standard states have been defined
for reporting heats of reactions
• Standard thermodynamic states are: 1 bar
pressure for all gases and 1 M concentration
for aqueous solutions
• A temperature of 25 °C (298 K) is often
specified as well
• The standard heat of reaction is the value
of the enthalpy change occurring under
standard conditions involving the actual
number of moles specified the the equation
coefficients
• An enthalpy change for standard conditions
is denoted H 
• For example, the thermochemical equation
for the production of ammonia from it
elements at standard conditions is:
N 2 ( g )  3H 2 ( g )  2NH3 ( g )
H   92.38 kJ
• The physical states are important
• The law of conservation of energy requires
2NH3 ( g )  N 2 ( g )  3H 2 ( g )
H   92.38 kJ
• Enthalpy is a state function
• An enthalpy diagram is a graphical
representation of alternate paths between
initial and final states
Two paths for the
formation of
carbon dioxide
gas. Each give the
same enthalpy
change.
Remember to include the physical states of reactants
and products in thermochemical equations.
•
•
•
Enthalpy changes for reactions can be
calculated by algebraic summation
This is called Hess’s Law: The value of
the enthalpy change for any reaction that
can be written in steps equals the sum of
the values of the enthalpy change of each
of the individual steps.
Enthalpy changes for a huge number of
reactions may be calculate using only a
few simple rules
• Rules for Manipulating Thermochemical
Equations:
1) When an equation is reversed the sign of the
enthalpy change must also be reversed.
2) Formulas canceled from both sides of an
equation must be for substances in identical
physical states.
3) If all the coefficients of an equation are
multiplied or divided by the same factor, the
value of the enthalpy change must likewise be
multiplied or divided by that factor.
• An enormous database of thermochemical
equations have been compiled:
– The standard heat of combustion is the amount
of heat released when 1 mol of a fuel
completely burns in pure oxygen gas with all
products brought to 25 °C and 1 bar
CH 4 ( g )  2O 2 ( g )  CO2 ( g )  2H 2 O(l ) H C  -890 kJ
Standard heats of combustion are always negative
and produce water in liquid form
– The standard enthalpy of formation of a
substance is the amount of heat absorbed when
1 mole of the substance if formed at 25 °C and
1 bar from its elements in their standard states
H 2 ( g )  12 O 2 ( g )  H 2 O(l ) H f  285.9 kJ/mol
• The standard enthalpy of formation for elements in
their standard states are zero
• These are the values most commonly used to
calculated standard enthalpy changes for reactions
• Standard enthalpies of formation are given in Table
7.2 and Appendix C
• Hess’s law can be restated in terms of
standard enthalpies of formation:






Sum
of

H
of
all
Sum
of

H

f
f of all
H reaction  


 of the products   of the reactants 
Example: Calculate the enthalpy of reaction for
2NO(g)+O2(g)2NO2(g)
ANALYSIS: Use Hess’s law and Table 7.2
SOLUTION:

H   2H fNO2  2H fNO  H fO2

 2(33.8 kJ) - (2(90.37 kJ)  0 kJ)  -113.1 kJ