Chemistry: Atoms First
Second Edition
Julia Burdge & Jason Overby
Chapter 10
Energy Changes in
Chemical Reactions
M. Stacey Thomson
Pasco-Hernando State College
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
10.1
Energy and Energy Changes
The system is a part of the universe that is of specific interest.
The surroundings constitute the rest of the universe outside the
system.
System
Surroundings
Universe = System + Surroundings
The system is usually defined as the substances involved in
chemical and physical changes.
Energy and Energy Changes
Thermochemistry is the study of heat (the transfer of thermal
energy) in chemical reactions.
Heat is the transfer of thermal energy.
Heat is either absorbed or released during a process.
heat
The SI energy unit is a Joule, J. Often calories are used and 1
calorie is the amount of heat required to raiseSurroundings
1 g of water by 1°C.
1 cal = 4.184 J
calorie is not the same as a nutritional calorie (Cal)
1 Cal=1000 cal
Energy and Energy Changes
An exothermic process occurs
when heat is transferred from the
system to the surroundings.
“Feels hot!”
2H2(g) + O2(g)
System
2H2O(l) + energy
heat
Surroundings
Universe = System + Surroundings
Energy and Energy Changes
An endothermic process occurs
when heat is transferred from the
surroundings to the system.
“Feels cold”
energy + 2HgO(s)
System
2Hg(l) + O2(g)
heat
Surroundings
Universe = System + Surroundings
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10.2
Introduction to Thermodynamics
Thermodynamics is the study of the interconversion of heat and
other kinds of energy.
In thermodynamics, there are three types
of systems:
An open system can exchange mass and
energy with the surroundings.
A closed system allows the transfer
of energy but not mass.
An isolated system does not exchange
either mass or energy with its
surroundings.
The First Law of Thermodynamics
The first law of thermodynamics states that energy can be converted
from one form to another, but cannot be created or destroyed.
ΔUsys + ΔUsurr = 0
ΔU is the change in the internal energy.
“sys” and “surr” denote system and surroundings, respectively.
ΔU = Uf – Ui; the difference in the energies of the initial and final
states.
ΔUsys = –ΔUsurr
Work and Heat
The overall change in the system’s internal energy is given by:
ΔU = q + w
q is heat
 q is positive for an endothermic process (heat absorbed by the
system)
 q is negative for an exothermic process (heat released by the
system)
w is work
 w is positive for work done on the system
 w is negative for work done by the system
Work and Heat
ΔU = q + w
Worked Example 10.1
Calculate the overall change in internal energy, ΔU, (in joules) for a system that
absorbs 188 J of heat and does 141 J of work on its surroundings.
Strategy Combine the two contributions to internal energy using ΔU = q + w
and the sign conventions for q and w.
Text Practice: 10.2 10.3 10.7 10.10 10.12
10.3
Enthalpy
The thermodynamic function of a system called enthalpy (H) is
defined by the equation:
H = U + PV
A note about SI units:
Pressure: pascal; 1Pa = 1 kg/(m . s2)
Volume:
cubic meters; m3
PV:
1kg/(m . s2) x m3 = 1(kg . m2)/s2 = 1 J
Enthalpy: joules
U, P, V, and H are all state functions.
Enthalpy and Enthalpy Changes
The enthalpy of reaction (ΔH) is the difference between the
enthalpies of the products and the enthalpies of the reactants:
ΔH = H(products) – H(reactants)
Assumes reactions in the lab occur at constant pressure
ΔH > 0 (positive) endothermic process
ΔH < 0 (negative) exothermic process
Thermochemical Equations
H2O(s)
H2O(l)
ΔH = +6.01 kJ/mol
Concepts to consider:
What is the system?
What are the surroundings?
ΔH > 0 endothermic
Thermochemical Equations
CH4(g) + 2O2(g)
CO2(g) + 2H2O(l)
ΔH = −890.4 kJ/mol
Concepts to consider:
What is the system?
What are the surroundings?
ΔH < 0 exothermic
Thermochemical Equations
Enthalpy is an extensive property.
Extensive properties are dependent on the amount of matter involved.
H2O(l) → H2O(g)
Double the amount of matter
2H2O(l) → 2H2O(g)
ΔH = +44 kJ/mol
Double the enthalpy
ΔH = +88 kJ/mol
Units refer to
mole of reaction as written
Thermochemical Equations
The following guidelines are useful when considering
thermochemical equations:
1) Always specify the physical states of reactants and products
because they help determine the actual enthapy changes.
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g)
different states
CH4(g) + 2O2(g)
CO2(g) + 2H2O(l)
ΔH = −802.4 kJ/mol
different enthalpies
ΔH = +890.4 kJ/mol
Thermochemical Equations
The following guidelines are useful when considering
thermochemical equations:
2) When multiplying an equation by a factor (n), multiply the ΔH
value by same factor.
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g)
ΔH = − 802.4 kJ/mol
2CH4(g) + 4O2(g)
2CO2(g) + 4H2O(g)
ΔH = − 1604.8 kJ/mol
3) Reversing an equation changes the sign but not the magnitude of
ΔH.
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g)
ΔH = − 802.4 kJ/mol
CO2(g) + 2H2O(g)
CH4(g) + 2O2(g)
ΔH = +802.4 kJ/mol
Worked Example 10.3
Given the thermochemical equation for photosynthesis,
6H2O(l) + 6CO2(g) → C6H12O6(s) + 6O2(g)
ΔH = +2803 kJ/mol
calculate the solar energy required to produce 75.0 g of C6H12O6.
10.4 Calorimetry
Calorimetry is the measurement of heat changes.
Heat changes are measured in a device called a calorimeter.
The specific heat (s) of a substance is the amount of heat required to
raise the temperature of 1 g of the substance by 1°C.
Specific Heat and Heat Capacity
The heat capacity (C) is the amount of heat required to raise the
temperature of an object by 1°C.
The “object” may be a given quantity of a particular substance.
heat capacity of 1 kg of water =
4.184 J
 1000 g = 4184 J/C
1 g  C
Specific heat capacity of water
Specific heat capacity has units of J/(g • °C)
Heat capacity has units of J/°C
heat capacity of
1 kg water
Specific Heat and Heat Capacity
The heat associated with a temperature change may be calculated:
q = msΔT
q = CΔT
m is the mass.
s is the specific heat.
ΔT is the change in temperature (ΔT = Tfinal – Tinitial).
C is the heat capacity.
Worked Example 10.4
Calculate the amount of heat (in kJ) required to heat 255 g of water from 25.2°C
to 90.5°C.
Calorimetry
A coffee-cup calorimeter may be used to measure the heat
exchange for a variety of reactions at constant pressure:
Calorimetry
Concepts to consider for coffee-cup calorimetry:
qP = ΔH
System:
reactants and products (the reaction)
Surroundings: water in the calorimeter
For an exothermic reaction:

the system loses heat

the surroundings gain (absorb) heat
qsys = −msΔT
qsurr = msΔT
qsys = −qsurr
The minus sign is used to keep
sign conventions consistent.
Worked Example 10.5
A metal pellet with a mass of 100.0 g, originally at 88.4°C, is dropped into 125
g of water originally at 25.1°C. The final temperature of both pellet and the
water is 31.3°C. Calculate the heat capacity C (in J/°C) of the pellet.
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Constant-Volume Calorimetry
Constant volume calorimetry is carried out in a device known as a
constant-volume bomb.
A constant-volume
calorimeter is an isolated
system.
Bomb calorimeters are
typically used to determine
heats of combustion.
qcal = −qrxn
Constant-Volume Calorimetry
To calculate qcal, the heat capacity of the calorimeter must be
known.
qcal = CcalΔT
qrxn = −qcal
qrxn = −CcalΔT
Worked Example 10.6
A chocolate chip cookie weighing 7.25 g is burned in a bomb calorimeter to
determine its energy content. The heat capacity of the calorimeter is 39.97
kJ/°C. During the combustion, the temperature of the water in the calorimeter
increases by 3.90°C. Calculate the energy content (in kJ/g) of the cookie.
Study Guide for Sections 10.1-10.4
DAY 26, Terms to know:
Sections 10.1-10.4 system, surroundings, thermochemistry, thermal energy, exothermic,
endothermic, open system, closed system, isolated system, first law of thermodynamics (law of
conservation of energy), enthalpy, calorimetry, specific heat, heat capacity, heat of reaction, coffee
cup calorimeter, bomb calorimeter
DAY 26, Specific outcomes and skills that may be tested on exam 4:
Sections 10.1-10.4
•Be able to describe potential energy changes for the system, kinetic energy and temperature
changes for the surroundings, and bond stability changes for the system in any endothermic or
exothermic system.
•Given either change in energy for the system or surroundings, be able to calculate the other
•Given two out of three of any of the following quantities, be able to calculate the third: change in
heat for the system, change in work for the system, overall change in energy for the system
•Given initial heat and final heat, be able to calculate the change in enthalpy for a system
•Be able to determine whether a process is endo or exothermic based on the mathematical sign for
the change in enthalpy
•Be able to calculate how the enthalpy change will be different if a chemical equation is reversed or
multiplied by a certain stoichiometric coefficient
•Given two of the following quantities, be able to calculate the third: temperature change, heat
change, heat capacity
•Given three of the following quantities, be able to calculate the fourth: temperature change, heat
change, mass, specific heat
•Be able to describe how to set up a coffee cup or bomb calorimeter to calculate changes in energy
based on measured changes in temperature and energy exchanges between system and
surroundings
Extra Practice Problems for Sections 10.1-10.4
Complete these problems outside of class until you are confident you have
learned the SKILLS in this section outlined on the study guide and we will
review some of them next class period. 10.5 10.7 10.11 10.13 10.23
10.27 VC10.3 VC10.8 10.31 10.35 10.37 10.41 10.97 10.115 10.121
Prep for Day 27
Must Watch videos:
https://www.youtube.com/watch?v=8m_FCe5aCqY (energy calculations, Tyler DeWitt)
https://www.youtube.com/watch?v=SV7U4yAXL5I (calculations with energy, crash course chemistry)
http://echem1a.cchem.berkeley.edu/#modules (Energy, UC-Berkeley, lessons 20.2-20.5 and 21.3-21.6)
http://www.youtube.com/watch?v=0mLFSDId6pQ (Bond enthalpy, TheChemistrySolution)
https://www.youtube.com/watch?v=9i26eK6xyNI (colligative properties, Isaacs)
https://www.youtube.com/watch?v=Y-wtTUPPeww (colligative properties II, Isaacs)
Other helpful videos:
http://ocw.mit.edu/courses/chemistry/5-111-principles-of-chemical-science-fall-2008/videolectures/lecture-16/ (MIT, start at 29 minute mark)
http://ps.uci.edu/content/general-chemistry-1b (UV-Irvine, lecture 6 and the first 30 minutes of lecture
7)
https://www.youtube.com/watch?v=7ugNp7-OTC8 (Bond enthalpy, Annette Tomory)
https://www.youtube.com/watch?v=CRIdLFod8uU (UC-Irvine)
Read Sections 10.5-10.8, 13.5-13.6