UNIT 3: Energy Changes and Rates of Reaction
Chapter 5: Energy Changes
Chapter 6: Rates of Reaction
UNIT 3 Chapter 5: Energy Changes
Chapter 5: Energy Changes
Energy changes are involved
in every chemical reaction—
even those that occur in our
bodies. While many
processes involve the release
of energy, others are
accompanied by the absorb
energy.
Sarah Reinertsen is a competitive
runner. She relies a great deal on
energy produced by the chemical
reactions in her cells.
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
5.1 The Nature of Energy and Heat
Chemical and physical changes in matter involve changes
in energy content. These changes in energy involve the
release or absorption of energy.
Thermochemistry is a branch of
chemistry that studies heat released
during chemical and physical processes.
Burning wood involves changes to the
matter that make up the wood. A change
in energy content also occurs. Energy is
released, mostly as heat and light.
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
Some Foundational Concepts for
Thermochemistry
Two categories of energy:
Kinetic Energy
Potential Energy
• energy of motion
• anything moving has
kinetic energy
• stored energy due to
the condition or
position of the object
SI unit of energy is the joule, J (1 kJ = 1000 J)
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
System and Surroundings
universe = system + surroundings
• the system is the sample being observed
• the surroundings is everything else
• interactions between a system and its surroundings
involve exchange of energy and matter
Three types of systems based on this type of exchange are:
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can exchange
energy and matter
with surroundings
can exchange energy, not
matter, with surroundings
cannot exchange matter or
energy with surroundings
UNIT 3 Chapter 5: Energy Changes
Section 5.1
Calculating the Amount of Heat
Entering and Leaving a System
Measurable properties of a system include:
volume, mass, pressure, temperature, and specific heat
capacity.
A calculated property of a system includes heat (Q) that
enters or leaves an object.
ΔT = Tfinal – Tinitial
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When heat enters a system, ΔT is positive and so is Q.
UNIT 3 Chapter 5: Energy Changes
Section 5.1
LEARNING CHECK
A 1.0 g sample of copper is heated from
25.0°C to 31.0°C. How much heat
did the sample absorb?
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Answer on
the next slide
UNIT 3 Chapter 5: Energy Changes
LEARNING CHECK
Q = mc∆T
Q = (1.0 g)(0.385 g/J°C)(6.0°C)
Q = 2.3 J
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Section 5.1
UNIT 3 Chapter 5: Energy Changes
Section 5.1
The First Law of Thermodynamics:
Energy is Conserved
The first law of thermodynamics states that:
energy can be converted from one form to another but
cannot be created or destroyed.
Since any change in energy of the universe must be zero,
ΔEuniverse = ΔEsystem + ΔEsurroundings = 0
ΔEsystem = –ΔEsurroundings
• if a system gains energy, that energy
comes from the surroundings
• if a system loses energy, that energy enters
the surroundings
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
Enthalpy
One way chemists express thermochemical changes is by a
variable called enthalpy, H. The change in enthalpy, ΔH,
of a system can be measured.
It depends only on the initial and final states of the system,
and is represented by
ΔH = ΔE + Δ(PV)
For reactions of solids and liquids in solutions, Δ(PV) = 0
If heat enters a system
If heat leaves a system
• ΔH is positive
• ΔH is negative
• the process is endothermic • the process is exothermic
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
The Second Law of Thermodynamics
The second law of thermodynamics states that:
when two objects are in thermal contact, heat is transferred
from the object at a higher temperature
to the object at the lower temperature until
both objects are the same temperature
(in thermal equilibrium)
When in thermal contact, energy from
hot particles will transfer to cold
particles until the energy is equally
distributed and thermal equilibrium is
reached.
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes:
Enthalpy of Solution
Three processes occur when a substance dissolves, each
with a ΔH value.
1. bonds between solute molecules or ions break
2. bonds between solvent molecules break
3. bonds between solvent molecules and solute molecules or
ions form
Sum of the enthalpy changes:
enthalpy of solution, ΔHsolution
The orange arrow shows
the overall ΔH.
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes:
Enthalpy of Phase Changes
Heat must be added to or removed from a substance in
order for the phase of the substance to change.
The ΔH for each phase change
has a particular symbol. For
example, ΔHmelt is called the
enthalpy of melting.
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The ΔH for one phase
change is the negative of
the ΔH for the opposite
phase change.
UNIT 3 Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes:
Chemical Changes
Every chemical reaction is associated with an
enthalpy of reaction, ΔHr.
For example:
The combustion of methane is
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
For every mole of methane combusted ΔHr = –890 kJ
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes:
Nuclear Changes
Alpha and Beta Decay:
Atoms of some elements emit alpha (α) or beta (β) particles,
transforming them into other elements with smaller masses.
Nuclear Fission:
• a heavier nucleus is split into smaller nuclei
• energy is released and the products are radioactive
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
Comparing Categories of Enthalpy Changes:
Nuclear Changes
Nuclear Fusion:
• two small nuclei combine to form a larger nucleus
• the most common reaction in the Sun is the fusion of
deuterium and tritium
In this nuclear fusion
reaction the highly
unstable helium-5
breaks down to helium4 and a neutron.
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UNIT 3 Chapter 5: Energy Changes
Section 5.1
Comparing Enthalpy Changes
• ΔH for physical changes, such as phase changes and
dissolving of a solute in solvent, are the smallest
• ΔHr are much larger than ΔH for physical changes, but
smaller than ΔH for nuclear changes.
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UNIT 3 Chapter 5: Energy Changes
Section 5.1 Review
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Section 5.1
UNIT 3 Chapter 5: Energy Changes
Section 5.2
5.2 Thermochemical Equations
and Calorimetry
Chemical reactions involve
• initial breaking of chemical bonds (endothermic)
• then formation of new bonds (exothermic)
ΔHr is the difference
between the total energy
required to break bonds
and the total energy
released when bonds form.
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UNIT 3 Chapter 5: Energy Changes
Section 5.2
Thermochemical Equations
Thermochemical equations include the enthalpy change.
Exothermic reaction:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + 890.8 kJ
Endothermic reaction:
N2(g) + 2O2(g) + 66.4 kJ → 2NO2(g)
The enthalpy term can also be written beside the equation.
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
N2(g) + 2O2(g) → 2NO2(g)
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ΔHr = –890.8 kJ
ΔHr = +66.4 kJ
UNIT 3 Chapter 5: Energy Changes
Section 5.2
Enthalpy Diagrams
Enthalpy diagrams clearly show the relative enthalpies of
reactants and products.
• For exothermic reactions, reactants have a larger enthalpy
than products and are drawn above the products.
• For endothermic reactions, the products have a larger
enthalpy and are drawn above the reactants.
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UNIT 3 Chapter 5: Energy Changes
Section 5.2
Molar Enthalpy of Combustion
Combustion reactions have their own symbol: ΔHcomb
For standard molar enthalpy of combustion data in
reference tables:
• reactants and products are at standard conditions
• refer only to the compound undergoing combustion
• are for combustion of 1 mol of the compound being
combusted, not for equations “as written.” For some
compounds to be written with a coefficient of 1, equations
are written with fractions as coefficients
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UNIT 3 Chapter 5: Energy Changes
Section 5.2
Reactant Amounts and Enthalpy of Reaction
• The enthalpy of a reaction is directly proportional to the
amount of substance that reacts.
• The combustion of 2 mol of propane releases twice as
much energy as the combustion of 1 mol of propane.
• Knowing the thermochemical equation, the enthalpy
change associated with a certain amount of reactants or
products can be calculated.
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UNIT 3 Chapter 5: Energy Changes
Section 5.2
LEARNING CHECK
Determine ΔHcomb for 15.0 g of propane.
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Answer on
the next slide
UNIT 3 Chapter 5: Energy Changes
Section 5.2
LEARNING CHECK
The molar mass of propane: 44.1 g/mol
Therefore, 15.0 g is 0.340 mol
ΔHcomb = ΔH°comb
ΔHcomb = (0.340 mol) (–2219.2 kJ/mol)
ΔHcomb = 755 kJ/mol
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UNIT 3 Chapter 5: Energy Changes
Section 5.2
Using Calorimetry To Study Energy Changes
Calorimeters are used to measure heat released or
absorbed by a chemical or physical process.
The water of a calorimeter is considered one system and the
container in which the process occurs another system.
These two systems are in thermal contact.
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For an endothermic process, heat will
be transferred from the water to the
process system. The temperature of
the water will decrease.
For an exothermic process, heat will
be transferred from the process to
the water. The temperature of the
water will increase.
UNIT 3 Chapter 5: Energy Changes
Section 5.2
Using a Simple Calorimeter
• nested polystyrene cups can be used (good insulators)
• a known mass of water is in the inner cup, where the
process occurs
• the process often involves compounds dissolved in water
• the change in temperature of the
water is measured; the solution
absorbs or releases energy
• the solution is dilute enough so
that the specific heat capacity
of water is used
A simple calorimeter.
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UNIT 3 Chapter 5: Energy Changes
Section 5.2
Using a Simple Calorimeter
If the “system” is the process being studied and the
“surroundings” is the water in the calorimeter:
The change in thermal energy caused by the process can be
calculated:
Q = mcΔT
m is the mass of the water, c is the specific heat capacity of
water, and ΔT is the change in temperature of the water
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UNIT 3 Chapter 5: Energy Changes
How to Use a Simple Calorimeter
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Section 5.2
UNIT 3 Chapter 5: Energy Changes
Section 5.2
Using Calorimetry Data To Determine the
Enthalpy of Reaction
Q = mwcwΔTw is the thermal energy absorbed or released by
the water.
The opposite sign of that value is the thermal energy released
or absorbed by the system (compounds in solution).
Since pressure is constant, Q exchanged by the system and
the water is the same as the enthalpy change of the system
ΔH. This is used to determine molar enthalpy change, ΔHr.
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UNIT 3 Chapter 5: Energy Changes
Section 5.2
Using Flame Calorimetry To Determine the
Enthalpy of Combustion
Flame calorimeters are flame-resistant
and often made of metals cans
A flame calorimeter
• is used for determining
ΔHcomb
• absorbs a great deal of
energy, which must be
included in energy
calculations
• is used for burning impure
materials like food; ΔHcomb
is reported in kJ/g
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UNIT 3 Chapter 5: Energy Changes
Section 5.2
Using Bomb Calorimetry To Measure
Enthalpy Changes during Combustion
A bomb calorimeter
• is used for more accurately
determining ΔHcomb
• determines ΔHcomb at
constant volume
• has a particular heat
capacity, C
Q = CΔT is used for bomb
calorimetry calculations
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Bomb calorimeters are much more
sophisticated than flame calorimeters
or simple calorimeters.
UNIT 3 Chapter 5: Energy Changes
Section 5.2
LEARNING CHECK
A chemical reaction is carried
out in a dilute aqueous solution
using a simple calorimeter.
Calculate the enthalpy change
of the reaction using the data
below.
mass of solution in calorimeter: 2.00 х 102 g
specific heat capacity of water: 4.19 J/g°C
Initial temperature: 15.0°C
Final temperature: 19.0°C
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Answer on
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UNIT 3 Chapter 5: Energy Changes
Section 5.2
LEARNING CHECK
Q = msolutioncsolution∆Tsolution
Q = (2.00 х 102 g)(4.19 J/g°C)(4.0°C)
Q = 3.35 х 103 J
This is the thermal change of the
“surroundings.”
∆Hsystem = –Q
∆Hsystem = –3.35 kJ
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UNIT 3 Chapter 5: Energy Changes
Section 5.2 Review
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Section 5.2
UNIT 3 Chapter 5: Energy Changes
Section 5.3
5.3 Hess’s Law
The enthalpy change of nearly any reaction can be
determined using collected data and Hess’s law.
The enthalpy change of any reaction can be determined if:
• the enthalpy changes of a set of reactions “add up to” the
overall reaction of interest
• standard enthalpy change, ΔH°, values are used
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UNIT 3 Chapter 5: Energy Changes
Section 5.3
Combining Sets of Chemical Equations
To find the enthalpy change for formation of SO3 from O2
and S8, you can use
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UNIT 3 Chapter 5: Energy Changes
Section 5.3
Techniques for Manipulating Equations
1.Reverse an equation
• the products become the reactants, and reactants become
the products
• the sign of the ΔH value must be changed
2.Multiply each coefficient
• all coefficients in an equation are multiplied by the same
integer or fraction
• the value of ΔH must also be multiplied by the same
number
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UNIT 3 Chapter 5: Energy Changes
Section 5.3
Standard Molar Enthalpies of Formation
Data that are especially useful for calculating standard
enthalpy changes:
standard molar enthalpy of formation, ΔH˚f
the change in enthalpy when 1 mol of a compound is
synthesized from its elements in their most stable form at
SATP conditions
• enthalpies of formation for elements in their most stable
state under SATP conditions are set at zero
• since formation equations are for 1 mol of compound,
many equations include fractions
• standard molar enthalpies of formation are in Appendix B
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UNIT 3 Chapter 5: Energy Changes
Section 5.3
Formation Reactions and Thermal Stability
The thermal stability of a substance is the ability of the
substance to resist decomposition when heated.
• decomposition is the reverse of formation
• the opposite sign of an enthalpy change of formation
for a compound is the enthalpy change for its
decomposition
• the greater the enthalpy change for the decomposition
of a substance, the greater the thermal stability of the
substance
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UNIT 3 Chapter 5: Energy Changes
Section 5.3
Using Enthalpies of Formation and Hess’s Law
For example:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
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UNIT 3 Chapter 5: Energy Changes
Section 5.3
LEARNING CHECK
Determine ∆H˚r for the
following reaction using the
enthalpies of formation that are
provided.
C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l)
∆H˚f of C2H5OH(l): –277.6 kJ/mol
∆H˚f of CO2(g): –393.5 kJ/mol
∆H˚f of H2O(l): –285.8 kJ/mol
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Answer on
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UNIT 3 Chapter 5: Energy Changes
Section 5.3
LEARNING CHECK
∆H˚r = [(2 mol)(∆H˚f CO2(g)) + (3 mol)(∆H˚f H2O(l))] –
[(1 mol)(∆H˚f C2H5OH(l)) + (3 mol)(∆H˚fO2(g)]
∆H˚r = [(2 mol)(–393.5 kJ/mol) + (3 mol)(–285.8 kJ/mol)] –
[(1 mol)(–277.6 kJ/mol) + (3 mol)(0 kJ/mol)]
∆H˚r = (–1644.4 kJ) – (–277.6 kJ)
∆H˚r = –1366.8 kJ
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UNIT 3 Chapter 5: Energy Changes
Section 5.3 Review
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Section 5.3
UNIT 3 Chapter 5: Energy Changes
Section 5.4
5.4 Energy Efficiency and
Energy Resources
Energy efficiency can be calculated using the equation:
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UNIT 3 Chapter 5: Energy Changes
Section 5.4
Using Energy Efficiently
A challenge in the development of energy efficient
technology is to find ways to best convert energy input into
useful forms.
For example, efficiency of appliances:
• conversion of input of electrical energy versus output of
energy usually all that is considered
• but should also consider efficiency
of the source of the electricity
Energy use distribution in Canadian
homes
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UNIT 3 Chapter 5: Energy Changes
Section 5.4
Conventional Energy Sources in Ontario
The three main sources of electrical energy in Ontario:
• nuclear power plants
• power plants that burn fossil fuels (natural gas and coal)
• hydroelectric generating stations
The distribution of energy sources in
Ontario
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UNIT 3 Chapter 5: Energy Changes
Section 5.4
Alternative Renewable Energy Sources
in Ontario
Renewable energy sources in Ontario:
• account for about 25% of energy production
• are projected to increase to as high as 40% by 2025
• include hydroelectric power (major source), wind energy
(currently ~ 1% and projected to 15% in 2025), and solar
energy (currently low but may be as high as 5% in 2025).
Much lower contributors are biomass, wave power, and
geothermal energy.
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UNIT 3 Chapter 5: Energy Changes
Section 5.4
What Is a “Clean” Fuel?
Different fuels have differing impacts on the environment.
One way this impact is measured is through emissions.
For example, CO2(g) emissions per kJ of energy produced.
Fuel
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kg CO2/ kJ energy
Anthracite coal
108.83
Oil
78.48
Natural gas
56.03
Nuclear
0.00
Renewables
0.00
UNIT 3 Chapter 5: Energy Changes
Section 5.4 Review
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Section 5.4