How are Energy and Chemistry Related?
Unit 13: Guided Notes
Main Idea:




Energy can change form and flow, but it is always conserved.
Thermochemical equations express the amount of heat released or absorbed
by chemical reactions.
The enthalpy change for a reaction can be calculated using Hess’s law.
Changes in enthalpy and entropy determine whether a process is
spontaneous.
New Skills:
 Calculate heat absorbed and heat released as temperature changes.
 Calculate calorimetry problems
 Determine enthalpy changes with phase changes
 Use Hess’s Law to determine enthalpy changes
 Calculate heat absorbed and heat released from reactions
 Determine entropy changes
 Calculate Free Energy
Academic Language:
Calorie
Calorimeter
Chemical potential energy
Energy
Enthalpy
Enthalpy of fusion
Enthalpy of vaporization
Entropy
Free energy
Heat
Hess’s Law
Joule
Law of conservation of energy
Second law of thermodynamics
Specific heat
Spontaneous reaction
Standard enthalpy of formation
surroundings
System
Thermochemical equation
Thermochemistry
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Unit 13 Homework:
CALM: http://calm.indiana.edu/
Book: Chapter 15
13.1 Energy and Heat
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CALM: 5 questions
p519 #1,2 , p521 #4,5
p525 #12-14
p528 #17, 18, 21
13.2 Thermochemical Equations
 CALM: 5 questions
 p532 #23-25
 p533 #27-30
13.3 Calculating Enthalpy Change
 CALM: 5 questions
 p537 #32-33
 p541 #35-37
13.4 Reaction Spontaneity
 CALM: 5 questions
 p548 #46-51
13.5 Accumulating Content and Skills
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How are Energy and Chemistry Related?
Chemistry Unit 13: Learning Goals and Objectives
13.1 Energy and Heat - Energy can change form and flow, but it is always conserved.

Define energy

Distinguish between potential and kinetic energy

Calculate the amount of heat absorbed or released by a substance as its
temperature changes.
13.2 Thermochemical Equations - Thermochemical equations express the amount
of heat released or absorbed by chemical reactions.

Explain the meaning of enthalpy and enthalpy change in chemical reactions
and processes.

Write thermochemical equations

Describe how energy is lost or gained during changes of state.

Calculate the heat absorbed or released in a chemical reaction.
13.3 Calculating Enthalpy Change – The enthalpy change for a reaction can be
calculated using Hess’s law.



Apply Hess’s law to calculate the enthalpy change for a reaction.
Calculate the ΔHrxn using thermochemical equations.
Determine the enthalpy change for a reaction using standard enthalpies of
formation.
13.4 Reaction Spontaneity - Changes in enthalpy and entropy determine whether
a process is spontaneous.

Define Entropy

Differentiate between spontaneous and nonspontaneous processes.

Explain how changes in entropy and free energy determine the spontaneity
of chemical reactions and other processes
13.5 Accumulating Content and Skills– Chemistry content is continuous and builds
on prior knowledge and skills. This section will combine this unit with previous units.
 Apply knowledge and skills from previous units to content learned in this
unit.
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13.1 Energy and Heat - Energy can change form and flow, but it is always
conserved.
Objective: Define energy
Energy
Energy is the ability to do work or produce heat. it exists in two basic forms:
1.
2.
Objective: Distinguish between potential and kinetic energy
Potential energy is energy due to the composition or position of an object.
 Example:

 Chemical Potential energy of the substance depends on its composition:

Kinetic energy is the related to the motion of an object or substance:
 Example:
 Kinetic energy of a substance is vibrational and random motion of its
representative particles.

Law of Conservation of Energy
The law of conservation of energy states that in any chemical reaction or physical
process, energy can be converted from on form to another, but it is neither created
nor destroyed.
 Also called the first law of thermodynamics.
 As the diver jumps from the diving board, all of the potential energy changes
to kinetic energy.

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Objective: Calculate the amount of heat absorbed or released by a substance as its
temperature changes.
Heat
Heat is the energy that is produced or transferred between objects.

 Energy transferred can be measured in different units:
 calorie (cal.) = the amount of energy to raise the temperature of 1
gram of water 1 °.
 joule (J) = SI unit for energy
 1 joule = .2390 calories, 1 cal = 4.184 joules
Specific Heat
The specific heat of any substance is the amount of heat required to raise the
temperature of one gram of that substance by one degree Celsius.
 Different substances have different specific heat values.
 Water = 4.184 J/g°C

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Heat Released and Heat Absorbed
Substances can absorb heat and release heat.


 q=smΔT
 q = heat (J), s = specific heat(J/gC), m= mass (g), and T= temperature
(C)
Practice Problem #1: In the construction of bridges and skyscrapers, gaps
must be left between adjoining steel beams to allow for the expansion and
contraction of the metal due to heating and cooling. The temperature of a
sample of iron with a mass of 10.0g changed from 50.4°C to 25.0°C with the
release of 114 J. What is the specific heat of iron?
Practice Problem #2: A metal rail released 250J of heat when its temperature
was lowered to 25.0°C . If the specific heat of the rail is 0.189 J/g°C , what
was the initial temperature of the rail?
Practice Problem #3: How much energy does it take to heat a 75,000 Liter
pool from 20. °C to 28°C ?
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Calorimetry
A calorimeter is an insulated device used for measuring the amount of heat
absorbed or released during a chemical or physical process.
Coffee Cup Calorimeter
 Constant pressure calorimeters or coffee cup calorimeters are often used
in school chemistry labs
Bomb Calorimeter
 Bomb calorimeters are constant volume calorimeters and are often used
for reactions involving gases.
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Calorimetry
When working a calorimetry problem, the heat transferred can be calculated by
measuring the the change in temperature in the calorimeter.
 -q (heat lost) = +q (heat gained)
 -smΔT=+smΔT
Practice Problem #4: Initially 125g of water in a calorimeter is 23.40°C. After
a 30.0 g metal sample at 115.0°C is placed into the calorimeter, the final
temperature of the water is 29.30°C. What is the specific heat of the metal
and how much energy was transferred?
13.2Thermochemical Equations- Thermochemical equations express the
amount of heat released or absorbed by chemical reactions.
Objective: Explain the meaning of enthalpy and enthalpy change in chemical reactions
and processes.
Thermochemistry
Thermochemistry is the study of heat changes that accompany chemical
reactions and phase changes. In thermochemistry we classify two parts of the
process.
 System =
 Surroundings =
 Universe =
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Enthalpy
Enthalpy (H) is the heat content and the total amount energy in a system.

 Therefore the change in heat of a system is represented, ΔH.
 Example: 4Fe(s) + 3O2(g) 2Fe2O3(s)
ΔH= -1625 kJ
Standard Enthalpy
Standard enthalpy, ΔH°, is the energy changes in a system when all reactants
and products are in their standard state.
 Standard state is the phase at 1 atm and 25°C (298K).

Objective: Write thermochemical equations
Thermochemical Equation
A thermochemical equation is a balanced chemical equation that includes
the physical states of all reactants and products and the energy change,
usually expressed as the change in enthalpy, ΔH.
Example: C6H12O6(s) + 6O2(g) 6CO2(g) + 6H2O(l)
ΔH= -2808kJ
 In this example,
Objective: Describe how energy is lost or gained during changes of state.
Heat of Fusion/Vaporization
The heat required to vaporize one mole of a liquid is called its (enthalpy) heat of
vaporization, (ΔHvap).
 H2O(l) H2O(g)
ΔHvap = 40.7 kJ/mol

The heat required to melt one mole of a solid substance is called its (enthalpy) heat
of fusion, (ΔHfus).
 H2O(s) H2O(l)
ΔHfus = 6.01 kJ/mol

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Heat of Vaporization and Fusion
Objective: Calculate the heat absorbed or released in a chemical reaction.
Practice Problem #5: A calorimeter is used to measure the heat evolved
when burning 1 mol of glucose.
C6H12O6(s) + 6O2(g) 6CO2(g) + 6H2O(l)
ΔH= -2808kJ
How much heat would be released if 54.0 g of glucose is burned?
Practice Problem #6: Which of the following processes is exothermic?
1. H2O(l)  H2O(g)
2. CH4(g)  CH4(l)
3. Br2(l)  Br2(s)
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13.3 Calculating Enthalpy ChangeObjective: Apply Hess’s law to calculate the enthalpy change for a reaction.
Hess’s Law
Hess’s law of heat summation states that if you can add two or more
thermochemical equations to produce an final equation, then the sum of the
enthalpy changes for the individual reactions is the enthalpy change for the final
reaction.
Hess’s Law Example:
 2S(s)+2Os(g)  2SO2(g)
ΔH = -594 kJ
 2SO2(g) + O2(g)  2SO3(g)
ΔH = -198 kJ
Objective: Calculate the ΔHrxn using thermochemical equations.
Hess’s Law Calculations
When combining reactions to for Hess’s law calculations, some reactions must be
altered.

 If the reaction is reversed, the sign of energy flow must be reversed
(enthalpy sign changes).
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Practice Problem #7: Use thermochemical equations a and b below to
determine ΔH for the decomposition of hydrogen peroxide (H2O2).
2H2O2(l)  2H2O(l) + O2(g)
a) 2H2(g) + O2(g)  2H2O(l)
ΔH=-572 kJ
b) H2(g) + O2(g)  H2O2(l)
ΔH=-188 kJ
Objective: Determine the enthalpy change for a reaction using standard enthalpies of
formation.
Enthalpy of Formation
Standard enthalpy (heat) of formation (ΔH°f) is defined as the change in
enthalpy that accompanies the formation of one mole of the compound in its
standard state from its elements in their standard states.
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S(s) + O
2

2(g)
® SO3(g)
ΔH°f = -396 kJ
Enthalpy of formation is based relative to a point of reference. The enthalpy of
formation of an element in its standard state is assumed to be 0.00kJ. Any
changes to that element in a chemical reaction has an energy change associated
with it.
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Enthalpies of Formation in Calculations
An example problem:
H2S(g) + 4F2(g)  2HF(g) + SF6(g)
ΔH°f =?
•
½H2(g) + ½F2(g) HF(g)
ΔH°f = -273 kJ
•
S(s) + 3F2(g)  SF6(g)
ΔH°f = -1220 kJ
•
H2(g) + S(s)  H2S(g)
ΔH°f = -21 kJ
Enthalpy Summation Equation
ΔH°rxn = ΣΔH°f(products) - ΣΔH°f(reactants)
Example problem:
H2S(g) + 4F2(g)  2HF(g) + SF6(g)
ΔH°f =?
•
½H2(g) + ½F2(g) HF(g)
ΔH°f = -273 kJ
•
S(s) + 3F2(g)  SF6(g)
ΔH°f = -1220 kJ
•
H2(g) + S(s)  H2S(g)
ΔH°f = -21 kJ
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Practice Problem #8: Use standard enthalpies of formation to calculate ΔH°rxn
for the combustion of methane.
CH4(g) + 2O2(g) CO2(g) + 2H2O(l)
13.4 Reaction Spontaneity- Changes in enthalpy and entropy determine
whether a process is spontaneous.
Objective: Define Entropy.
Entropy
Entropy (S) is a measure of the number of possible ways that the energy of a
system can be distributed.
 This is related to the “positional probability” or the freedom of the
system’s particles to move and the number of ways they can be arranged.
 Solid –
 Gas –
Second Law of Thermodynamics
The second law of thermodynamics states that spontaneous processes always
proceed in such a way that the entropy of the universe increases.
 Entropy is sometimes considered to be a measure of the disorder and
randomness of the particles that make up a system.
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Changes in Entropy
Changes in entropy are calculated in a similar way that enthalpy is:
ΔSsystem = Sproducts – Sreactants

 If the entropy of a system decreases during a reaction, ΔSsystem is negative.
Sproducts < Sreactants
1. Entropy changes associated with changes in state can be predicted:
a. H2O(l)  H2O(g)
ΔSsystem
2. The dissolving of a gas in a solvent always results in a decrease in entropy
a. O2(g) O2(aq)
ΔSsystem
3. Assuming no change in physical state occurs, the entropy of a system usually
increases when total number of gas molecules increases.
a. 2SO3(g) 2SO2(g) + O2(g)
ΔSsystem
4. With a few exceptions, entropy increases when a solid or a liquid dissolves in
a solvent.
a. NaCl(s)  Na+(aq) + Cl-(aq)
ΔSsystem
5. The random motion of the particles of a substance increases as its
temperature increases.
a. When temperature increase, ΔSsystem
Practice Problem #9: What is the sign of ΔSsystem for each of the following
changes:
1. CH3OH(l)  CH3OH(g)
2. CaCO3(s) CaO(s) + CO2(g)
3. CuS(s) + 2O2(g)  CuSO4(s)
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Objective: Differentiate between spontaneous and nonspontaneous processes.
Spontaneous Process
A reaction that occurs with no outside intervention is a spontaneous process.
 Some spontaneous processes require some outside energy to get the
process started.

 example: a car rusts spontaneously.
Spontaneity is dependent on two things:
 The change in enthalpy of a reaction

 Rust: 4Fe(s) + 3O2(g)  2Fe2O3(s)
ΔH= -1625 kJ
 The change in entropy of a reaction

Objective: Explain how changes in entropy and free energy determine the spontaneity
of chemical reactions and other processes
Free Energy
Willard Gibbs, a Yale physicist, combined the idea of enthalpy and entropy in
order to determine if a reaction is spontaneous. He called the equation Gibbs
Free Energy:
ΔGsystem = ΔHsystem – TΔSsystem
 ΔH is enthalpy change, T is temperature in Kelvin, and ΔS is entropy
change
 If ΔGsystem
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Practice Problem #10: Determine the reaction spontaneity for the following
reaction at 382K:
ΔHsystem =145 kJ and ΔSsystem = 322 kJ
13.5 Accumulating Content and Skills– Chemistry content is continuous and
builds on prior knowledge and skills. This section will combine this unit with previous
units.
 Apply knowledge and skills from previous units to content learned in this
unit.
o Predict the typical enthalpy and entropy changes for the four main
types of reactions:
 Synthesis

Decomposition

Combustion

Replacement
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