Chapter 5 - Thermochemistry
Being able to measure the amount of energy that flows
either into or out of a system is an important part of
chemistry. Chemists can measure the amount of energy
released by burning fuel. Food also contains energy that can
be measured in calories. How can one figure out how much
potential energy a particular type of food contains, or
chemical reaction may have?
Thermochemistry
• Energy
• Is defined as the ability to do work.
• Work = (force Applied) x (distance object move)
W=F.d
Thermochemistry
• Energy
• Energy can take on different forms;
• Heat
• Light
• Sound
Thermochemistry
• Energy
• Kinetic Energy (Latin: Kinesis = to move)
• The energy of motion.
• If an object, molecule,
or atom moves as a
result of a force being
applied, we can say
that it possesses kinetic
energy.
Thermochemistry
• Energy
• Kinetic Energy (Latin: Kinesis = to move)
• In chemistry, molecules that move possess
kinetic energy.
• Kinetic energy
creates heat.
• Heat is a measure
of kinetic energy.
Thermochemistry
• Energy
• Kinetic Energy
KE = 1/2m.v2
m = mass (kg)
v = velocity (m/s)
Thermochemistry
• Energy
• Potential Energy
• A stick of dynamite has potential energy.
• When it explodes, the potential energy is
converted into kinetic energy (heat).

Thermochemistry
• Energy
• Potential Energy
• The amount of potential energy contained in
food is a popular topic.
• Some foods contain
a very large amount
of kinetic energy.
Thermochemistry
• Energy
• Potential Energy
• The amount of potential energy contained in
food is expressed in Calories.
• There are 813 Calories in a Cinnabon.
• There is enough energy in 1 Cinnabon to
heat 813000 grams of water by
1°C. That’s 813 L (215 gallons!)
Thermochemistry
• Energy
• Potential Energy
Eel = kQ1Q2
d
The electrostatic potential energy is equal to
the product of two electrical charges times a
proportionality constant (k = 8.99 x 109 J.m/C2)
and inversely proportional to the distance
separating them, d (meters).
Thermochemistry
• Energy
• Units for Measuring Energy
• The most common units for energy is the
Joule (J), calorie (cal), and the Calorie (Cal).
• 1 calorie is defined as the amount of energy
required to raise the temperature of 1.o
gram of water by 1.0°C.
Thermochemistry
• Energy
• Units for Measuring Energy
• Therefore 1 Calorie is the amount of energy
required to raise 1.0 kg of water by 1.0°C.
1 calorie = 4.18 Joules
1000 calories = 1 Calorie
Thermochemistry
• Energy
• Units for Measuring Energy
• Convert 2.0 x 103 calories into Joules.
• 4.5 x 104 J = ? Cal
Thermochemistry
• Energy
• Chemical Reactions and Energy
• A chemical reaction that releases energy is
called an exothermic reaction.
• The prefix ‘exo’ means
outside, or to leave.
Thermochemistry
• Energy
• Chemical Reactions and Energy
• An Endothermic reaction absorbs energy
from its surroundings.
• The prefix ‘endo’ means
to enter or absorb.
• An ice-pack absorbs
energy from your body
in the form of body heat.
Thermochemistry
• Energy
• Physical Changes and Energy
• Condensation and freezing are examples of
an exothermic process.
• Energy has to be released when a substance
changes from a gas to a liquid, or a liquid to
a solid.
Thermochemistry
• Energy
• Physical Changes and Energy
• Vaporization, evaporation, and melting are
examples of an endothermic phase change.
• Energy is absorbed when a substance
changes from a liquid to a gas or a solid to a
liquid.
Thermochemistry
• Energy
• Heat Capacity (Specific Heat)
• The amount of energy required to raise the
temperature of a substance by 1°C.
• The Specific Heat Capacity (Specific Heat) of
a substance is the amount of energy
required to raise the temperature of 1.0
gram of the substance by 1.0 °C.
Thermochemistry
• Energy
• Heat Capacity (Specific Heat)
• The unit of the calorie is defined using the
specific heat of water.
• 1 calorie is the defined as the amount of
energy required to raise the temperature of
1.0 gram by 1.0°C.
Thermochemistry
• Energy
• Energy Calculations
• We can relate energy, mass, and temperature
with the following equation;
q = m x c x ΔT
q = energy (J. cal, or Cal)
m = mass (g or Kg)
c = specific heat constant (J/g.°C)
The units of specific heat may vary.
(J/g.°C) (cal/g.°C) (cal/kg.°C) (Cal/g.°C) (Cal/kg.K)
Thermochemistry
• Energy
• Specific Heat Values
Thermochemistry
• Energy
• The sign of ‘q’
• If ‘q’ has a positive value, energy is being
absorbed by the substance.
• If ‘q’ has a negative value, energy is being
released from the substance.
Thermochemistry
• Energy Problems
• How many Joules of energy is required to
heat 1.0 L of water from 23.0°C to
100.0°C? (cH2O = 4.18 J/g.°C)
Thermochemistry
• Energy Problems
• How many calories of energy is lost when
a 400.0 gram piece of iron cools from
200.0°C to 50.0°C? (cFe = 0.11 cal/g.°C)
Thermochemistry
• Calorimetry
• A calorimeter is an insulated container
that can measure the amount of energy
that either flows into
or out of an object or
chemical reaction.
• Water is usually used
to capture, or release
the energy.
Thermochemistry
• Calorimetry
• When using a calorimeter, it is assumed that
the energy lost or gained by the object or
reaction, is equal to the energy lost or gained
by the water in the calorimeter.
q (lost by object) = q (gained by calorimeter)
or
q (gained by object) = q (lost by calorimeter)
Thermochemistry
• Calorimetry
• For an exothermic reaction, we can assume
that;
-q (reaction or object) = +q (calorimeter)
-mcΔT (reaction or object) = +mcΔT (calorimeter)
Thermochemistry
• Calorimetry
• A 200.0°C piece of lead, with a mass of 12.0 g,
is placed into a calorimeter that holds 300.0
grams of water. The temperature of the water
rose by 18.0°C. Calculate the experimental
value for the specific heat constant for lead.
(specific heat of water = 4.18 J/g . °C)
Thermochemistry
• Calorimetry
• 10.0 mL of a 1.0 M aqueous solution of HNO3 is
added to 10.0 mL of a 1.0 M aqeous solution of
NaOH. Each solution is recorded as having a
temperature of 23.0°C before they are mixed.
After the two solutions are mixed, the
temperature increases to 30.0°C . Assuming
that each solution is mostly water, calculate the
enthalpy change (ΔH) of this chemical reaction.
Thermochemistry
• Energy and Phase Changes
• The graph shows the continuous heating of
water versus time.
• Describe what
is happening.
Thermochemistry
• Energy and Phase Changes
• The temperature of a substance remains
constant when it is going through a phase
change.
• We can calculate
the amount of
energy absorbed in
the sloped regions
by using q=mcΔT.
Thermochemistry
• Energy and Phase Changes
• We can calculate the amount of energy
absorbed during the phase changes by using
the heat of fusion,
ΔHfus, or the heat
of vaporization,
ΔHvap.
Thermochemistry
• Energy and Phase Changes
• Calculate the amount of energy absorbed when
5.00 grams of ice melts and changes into water.
ΔHfus H2O = 334 J/g
Thermochemistry
• Energy and Phase Changes
• Calculate the amount of energy released when
18.00 grams of steam condensed and changes
into water.
ΔHvap H2O = 2260 J/g
Thermochemistry
• Energy and Phase Changes
• Calculate the amount of energy that is required
to heat water from -20.0°C to 80.0°C.
ΔHfus H2O = 334 J/g
Thermochemistry
• Energy and Phase Changes
• Cooling Curves
Thermochemistry
• Thermochemical Equations
• The change in the amount of energy of a
chemical equation is expressed as an
Enthalpy Change (ΔH).
• Endothermic Reaction = +ΔH
• Exothermic Reaction = -ΔH
Thermochemistry
• Thermochemical Equations
• We can calculate the theoretical ΔH of a
chemical reaction if we know the Heats of
Formation (ΔHf) of the reactants and
products.
N2(g) + 2H2(g)  2NH3(g)
ΔHf N2 = 0 kJ/mol
ΔHf H2 = 0 kJ/mol
ΔHf NH3 = -46 kJ/mol
Thermochemistry
• Thermochemical Equations
• Hess’s Law – The Heat of a chemical
reaction (ΔHreaction) can be calculated by
taking the difference of the heat of
formation of the product(s) (ΔHproducts) and
heat of formation of the reactant(s)
(ΔHreactants).
• ΣΔHreaction = ΣΔHproducts - ΣΔHreactants
Thermochemistry
• Thermochemical Equations
Thermochemistry
• Thermochemical Equations
• Hess’ Law
• Calculate ΔHreaction for the combustion of
methane, CH4(g) using Hess’ Law.
Thermochemistry
• Thermochemical Equations
• Hess’ Law
• Calculate ΔHreaction for the neutralization
reaction between HCl and NaOH.