Chapter 12 “Heat in
Chemical Reactions”
Where did all of that
energy come from?
T. Witherup 2006
OBJECTIVES
• Distinguish between “exothermic” and
“endothermic” chemical reactions.
• Explain what is meant by enthalpy and
enthalpy change.
• Define standard enthalpy change and
explain its use.
• State and apply “Hess’s Law” to determine
ΔH for a reaction.
• Describe calorimetry and how it is used to
determine heats of reactions.
• Compare “heat” and “temperature.”
2-1 Chemical Reactions That
Involve Heat
• What do you visualize when thinking about
chemical reactions? You probably imagine…
• Heat, Light, Sound
• These are energy, and chemical reactions involve
energy because chemical bonds are being broken
and made!
• Energy: the capacity to do work.
• What kinds of energy are there?
•
•
•
•
•
Radiant (solar)
Thermal (motion of atoms & molecules)
Chemical (stored within chemical bonds)
Potential (available because of an object’s position)
Kinetic (energy of motion)
12-1 Thermochemistry &
Thermodynamics
• Heat is the transfer of thermal
energy between two objects at
different temperature.
• Thermochemistry is the study of
heat changes that occur in chemical
reactions.
• ‘Thermes’ is Greek for “heat.”
• Thermodynamics is the broader
study of energy and work changes
during such processes.
12-1 Exothermic vs.
Endothermic Reactions
• Exothermic reactions release heat energy.
• ‘Exo’ means outside.
• Example: Burning natural gas (methane, CH4) to
keep warm in the winter.
• CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + Heat
• Notice that “heat” is a product!
• Endothermic reactions absorb heat energy.
• ‘Endo’ means inside.
• Example: Decomposing HgO into mercury and
oxygen. (Lavoisier used this reaction to discover
oxygen.)
• 2HgO(s) + Heat  2Hg(l) + O2 (g)
• Now note that “heat” is a reactant!
12-1 Exothermic vs. Endothermic
Reactions (Examples)
• U.S. Army “MREs” (“Meals ready to eat”)
• Magnesium reacts with water to release heat and hydrogen
gas. Write the reaction.
• Mg + 2H2O  Mg(OH)2 + H2 + Heat
• MREs heat the food to 60° C in 15 minutes!
• EXOTHERMIC Reaction
• Ammonium Nitrate (A common fertilizer)
• NH4NO3 dissolves in water, causing a drop in temperature.
Write the reaction.
• NH4NO3 + H2O + Heat  NH41+ (aq) + NO31- (aq)
• NH4NO3 is used in cold packs for sports injuries.
• ENDOTHERMIC Reaction
• Note: NH4NO3 can also undergo an EXOTHERMIC reaction
when it explodes, hence its use as a blasting agent in
mining and (unfortunately) in terrorists’ weapons, such as
that used in the Oklahoma City attack.
12-2 Heat and Enthalpy
Changes
• Remember the “Law of Conservation of
Energy”?
• It is also the First Law of
Thermodynamics:
• “Energy cannot be created or destroyed, but may
be converted from one form to another.”
• To study thermochemistry we must be
careful about the ‘system’ involved.
• Open System: allows exchange of both energy
and mass.
• Closed System: allows the exchange of energy,
but not mass.
• Isolated System: allows the exchange of neither
energy nor mass.
• See diagrams on next slide.
Types of “Systems”
OPEN
Mass Loss
or Gain
CLOSED
ISOLATED
Energy Loss
or Gain
A New Term: Enthalpy
• Most chemical processes we use occur at
normal, constant atmospheric pressure, and heat
energy is either gained or lost during the
chemical reaction.
• Chemists use a special term, enthalpy, to
describe heat changes for chemical reactions at
constant pressure.
• Enthalpy is the energy of a substance plus a
small adjustment for pressure-volume work.
• Enthalpy (H) = E + PV
= Internal Energy + Pressure-Volume (Work)
= Heat absorbed or released by a reaction
• “Enthalpy” is from German (enthalpein, to warm).
Enthalpy Change (ΔH)
• Enthalpy change, ΔH, for a chemical reaction that
occurs at constant pressure is the heat released or
absorbed in the reaction.
• For the reaction
Reactants  Products
ΔH is just the difference between the enthalpy of
the products minus the enthalpy of the reactants:
ΔH = Hproducts – Hreactants
(Here the Δ means “change” or “a difference.”)
Think of enthalpy in terms of a savings account
balance. You are just keeping track of money going
into the account and the money spent vs. the net
balance.
Some Conventions (Rules)
• We use diagrams to show enthalpy
changes.
• For exothermic reactions, Hproducts is
lower than Hreactants so ΔH is negative.
• Heat is released (lost by the system).
• For endothermic reactions, Hproducts is
higher than Hreactants so ΔH is positive.
• Heat is absorbed (gained by the system).
• See diagrams on the next slide.
Exothermic Reaction Diagram
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + Heat
CH4(g) + 2O2(g)
E
N
T
H
A
L
P
Y
ΔH = -890kJ
(EXOTHERMIC; HEAT
RELEASED FROM THE
‘SYSTEM’ TO THE
SURROUNDINGS)
CO2(g) + 2H2O(g)
REACTION PROGRESS (TIME)
NOTE: Hproducts IS
LOWER THAN
Hreactants
Endothermic Reaction Diagram
2HgO(s) + Heat  2Hg(l) + O2 (g)
2Hg(l) + O2(g)
E
N
T
H
A
L
P
Y
2HgO(s)
ΔH = +181kJ
(ENDOTHERMIC; HEAT
ABSORBED BY THE
‘SYSTEM’ FROM THE
SURROUNDINGS)
REACTION PROGRESS (TIME)
NOTE: Hproducts IS
HIGHER THAN
Hreactants
“Standard States”
• To be exact when comparing enthalpy
changes chemists must define the conditions,
such as pressure, temperature and phase.
• The “Standard State” of a substance is its
pure form at the standard pressure of 1
atmosphere and 25° C.
• Standard state of oxygen is as the diatomic gas.
• Standard state of carbon is graphite.
• Standard state of phosphorus is P4 (white
phosphorus).
• Enthalpy changes measured under “standard
state” conditions are called “Standard
Enthalpy Changes” (ΔH°).
Using Enthalpy Changes
How much heat will be released when 6.440g
of sulfur reacts with excess oxygen by the
reaction:
2S (s) + 3O2 (g)  2SO3 (g)
Solution:
6.440g S X
1 mol S
32.07g S
X
ΔH° = -791.4kJ
-791.4kJ
2 mol S
=
•See sample problems (p. 386) & 12-2 worksheets.
-79.46 kJ
12-3 Hess’s Law
• Consider the energy changes that occur as two
groups of adventurers climb a mountain using
different trails.
• Group 1 makes the difficult climb straight up the face of
the mountain.
• Group 2 takes the gently sloping trail that winds slowly
up the mountain.
• What is their relative Potential Energy (PE) at the base
camp and at the summit?
• Hess’s Law: If a series of chemical reactions
are added together, the enthalpy change for the
net reaction will be the sum of the enthalpy
changes for the individual steps.
Hess’s Law Example
• Calculate the standard enthalpy change for
the reaction 2C(s) + O2(g)  2CO(g) from
• C(s) + O2(g)  CO2(g) ΔH = -393.5 kJ
• 2CO(g) + O2(g)  2CO2(g) ΔH = -566.0 kJ
• (Ans.: -221.0 kJ)
• See also sample problems (p390-391) & 123PP sheet.
• The trick is to simply combine equations so
that only the relevant parts are present
(everything else cancels out). Remember,
whatever you do to the equations, you must
do to the value of ΔH for the reaction.
12-4 Calorimetry
• What happens during chemical reactions?
• During an exothermic reaction, the surroundings
gain heat from the ‘system.’
• During an endothermic reaction, the surroundings
lose heat to the ‘system.’
• Calorimetry is the study of heat flow and
heat measurement during chemical
reactions.
• Calorimetry experiments use accurate
measurements of temperature changes to
determine the enthalpy changes (heat flow).
Heat vs. Temperature
• Heat Capacity: the amount of heat needed to raise
the temperature of an object by 1 degree Celsius
(1° C).
• Specific Heat: the heat capacity of one gram of a
substance.
• Specific Heat of water = 4.184 J/g.C°
• To raise the temperature of one gram of liquid water
by 1 degree C requires 4.184 J of heat energy, or 1
calorie (cal.)
• Recall: 1,000 calories = 1 kilocalorie (1 kcal) = 1
Calorie (the ‘food’ unit)
• Note that heat and temperature are related, but
different, concepts.
• We detect heat changes by measuring temperature
changes.
• A small temperature change does not mean a small
amount of heat transfers! (Consider melting a lake!)
• Use specific heat and mass to measure heat changes.
A Calorimetry Experiment
•
Calorimeter: a well insulated container filled with
a known mass of water, a way to conduct a
chemical reaction, a stirrer and a thermometer.
• Measure the initial temperature of the water (Ti).
• Conduct the chemical reaction.
• Measure the temperature periodically.
• Determine the final temperature (Tf).
• Note that all heat released by a reaction is gained
by the water in the calorimeter, and vice versa!
• Mathematically we may say:
(qrxn = -qsur)
• Calculate the quantity of heat absorbed by the water (qsur) from the
mass of water, the Specific Heat of water and the temperature
change (Tf – Ti). (And vice versa.)
qsur = m X C X (Tf – Ti)
• The sign of
qsur indicates the direction of heat flow for the reaction.
• Exothermic reaction, heat goes to the water: (-) sign.
• Endothermic reaction, heat goes from the water: (+) sign.
Calorimetry Example
•
When 1.008g of hydrogen gas reacts with chlorine gas in a calorimeter
containing 500.00g of water, the temperature rises from 25.00°C to
29.24°C. Calculate ΔH for the process H2(g) + Cl2(g)  2HCl(g) ΔH = ?
• Use qsur = m X C X (Tf – Ti)
• qsur = (500.00g) X (4.184J/g.C°) X (29.24 - 25.00 C°) = 8870J
• qsur = 8870J = -qrxn
(An exothermic reaction.)
• Therefore qrxn = -8870J.
• 1.008g H2 X (1 mol H2/2.00g H2) = 0.504 mol H2, which
released this -8870J of energy.
• So for the overall reaction:
(-8870J/0.502 mol H2) = ( -17,700J/1mol H2)
Or -17.7kJ/mol H2
Fuel for Our Body
• Our body is like a chemical factory, taking in
raw materials of various kinds to make new
products that keep us alive and help us grow.
• These reactions involve bond-breaking and
bond-making, and they meet the energy needs
of our body.
• Carbohydrates and fats are major food sources,
and they release lots of energy (exothermic
reactions).
• Glucose releases -2803kJ/mol C6H12O6
• See Table on page 399.
• “Calories Count” Project
• Recall 1000 cal = 1 Kcal = 1 Cal (the food kind!)
12-5 What is Heat?
• “Caloric Theory”
• Heat was thought of as an invisible, weightless
fluid capable of flowing from a hot object to a
colder one.
• Benjamin Thompson (Count Rumford)
showed that friction heat between objects
was continuous as long as two objects
were rubbed together.
• This was not consistent with ‘Caloric Theory.’
• James Joule used a paddlewheel
experiment to convert mechanical energy
into heat, thus showing there is a
‘mechanical equivalent of heat.’
• Modern “Kinetic Theory”
• Heat is defined as the transfer of kinetic energy
from a hotter object to a colder one.
OBJECTIVES
• Distinguish between “exothermic” and
“endothermic” chemical reactions.
• Explain what is meant by enthalpy and
enthalpy change.
• Define standard enthalpy change and
explain its use.
• State and apply “Hess’s Law” to determine
ΔH for a reaction.
• Describe calorimetry and how it is used to
determine heats of reactions.
• Compare “heat” and “temperature.”