CH 6: Thermochemistry
6.1 Nature of Energy
• Thermochemistry – study of energy
changes during chemical reactions
– Aspects of thermochemistry are studied in
both physics and chemistry
6.1
• Energy – the capacity to do work or
produce heat
• Law of conservation of energy states that
energy can be converted from one form to
another, but it cannot be created or
destroyed.
– Energy of the universe is constant!
6.1
Two forms of energy
1. Potential energy – stored energy
• Energy of position or composition
• Examples
2. Kinetic energy – energy of motion
• Heat, light, electricity
• In this chapter our focus will (eventually) be on the
heat aspects of thermochemistry.
6.1
• Consider the diagram on page 237.
• Ball A has potential energy due to its
position
– Some of this energy is released as heat as A
rolls down the hill
– Ball A hits Ball B and work is done as B
moves up the incline
• Work = force acting over a distance
6.1
• The original potential energy of A is equal
to the new potential energy of B plus the
heat released as friction.
• Energy can be released as:
– Heat (friction)
– Work (ball rolling)
6.1
• The total energy released when ball A rolls
down the hill is fixed – how it’s released is
not.
– How much is released as heat vs. work
depends upon the conditions
– Total energy released does not depend on the
pathway.
6.1
• Total energy change is a state function.
– State function is a property of a system that
depends only on its current state not its past
• See last paragraph on page 237
– Energy of a system is a state function.
• Heat and work are not state functions – they
depend on the path taken
6.1 Chemical Energy
• When studying chemical change we
consider the system and the surroundings.
– System includes the reactants and products
of a given reaction
– Surroundings are everything else in the
universe!
• Including the water if the reaction is done in
solution
Chemical Energy
• Exothermic reactions – energy flows out of
the system
– Products are of lower potential energy than
the reactants
• Energy is released to the surroundings
• Energy lost by the system = energy gained by the
surroundings
– There’s a loss of energy by the system….
D Energy of the system is negative
Chemical Energy
• Endothermic reactions – energy flows into
the system
– Products are of higher potential energy than
the reactants
• The system absorbs energy from the surroundings
• The energy gained by the system = the energy lost
by the surroundings
There’s a gain of energy by the system…….
D Energy of the system is positive
Chemical Energy
• Exothermic: D E < 0
– System _______ energy
• Endothermic: D E > 0
– System ________ energy
More on Energy of the System
• The energy of the system can change as a
result of 2 factors:
– Heat (q)
– Work (w)
• Change of energy of the system = heat
flowing in/out of system + work being done
to/by the system
DE= q + w
D E of the System
• Heat (q)
– When heat flows into the system, q > 0
• Endothermic
– When heat flows out of the system, q < 0
• Exothermic
• Work (w)
– When work is done to the system, w > 0
– When work is done by the system, w < 0
More on Work
• Most common form of work done by a
system is the expanding or compressing of
a gas, called pressure – volume work
– Work – force applied over a distance
• Moving an object a distance = work
Pressure Volume Work
• Consider a gas in a cylinder with a movable piston on
top
• Work = force x distance
Force = pressure x area
Distance = change in height of gas in cylinder (D h)
• Work = P x A x D h
D Volume
• Page 241
Pressure Volume Work
• Work = P x D V
– The sign ( +/-) on work is assigned so that:
• When the gas expands, it is doing work on the
surroundings (w < 0)
• When the gas is compressed, work is done on the
system (w > 0)
Pressure Volume Work
• Final version of the equation that shows
both magnitude and sign on work:
W=-PxDV
• When the gas expands D V is positive and
w < 0 (work is done by the system)
• When the gas is compressed D V is negative
and w > 0 (work is done on the system)
6.2 Entahalpy
• Enthalpy (H)
– Enthalpy is defined as: H = E + PV
• E is the energy of the system
• P is the pressure of the system
• V is the volume of the system
Enthalpy
• In chemistry we consider enthalpy at
constant pressure
– After much math this results in the formula:
D H = qp
D H is called the heat of the reaction
DH is a measure of the flow of heat (q)
into/out of the system at constant pressure
Enthalpy
• When heat leaves the system D H < 0
– Exothermic process
• When heat enters the system D H > 0
– Endothermic process
Finally – the Applications!
CH4 + 2 O2

CO2 + 2 H20 + energy
D H = - 890 kJ
1. Is the reaction exothermic or endothermic?
2. How much energy will be ___________ if 6.50
grams of CH4 is burned at constant pressure?
Next…..#44 on page 277
Calorimetry
• The heat changes associated with a
chemical reaction are often measured in a
calorimeter.
Calorimetry
• Exothermic: The heat released by the
reaction is used to heat up a known
quantity of water.
• More heat released the hotter the water gets
• Endothermic: The heat absorbed by the
reaction comes from the water
• More heat absorbed the colder the water gets
Terms
• Terms all describe the energy needed to
heat or cool some amount of a given
substance
– Heat Capacity (C)
– Specific heat capacity
– Molar heat capacity
Terms
• Heat Capacity (C)
– Amount of energy needed to raise the
temperature of a substance by 10C
– Units: J/0C
– Substance is an entire/specific object…e.g.
Terms
• Specific Heat Capacity
– Amount of energy needed to raise the
temperature of 1 gram of a substance by 10C
– Units: J/g0C
– See page 245
Terms
• Molar Heat Capacity
– Amount of energy needed to raise the
temperature of 1 mole of a substance by 10C
– Units: J/mol 0C
Molar heat capacity = specific heat x molar mass
J/mol 0C =
J/g 0C
x
g/mol
Calculating Heat Capacities
Heat Capacity =
Specific Heat =
Page 277: 52, 54
heat absorbed
DT
energy
(mass) (D T)
Determining Specific Heat of a
Metal
• Lab demonstration of experiment 27:E,
page 347 of the lab manual
– Assume the heat capacity of the ccc is 0
J/g0C