George Mason University

1
George Mason University General Chemistry
211 Chapter 6 Thermochemistry Energy Flow and
Chemical Change Acknowledgements Course Text
Chemistry the Molecular Nature of Matter and
Change, 6th edition, 2011, Martin S. Silberberg,
McGraw-Hill The Chemistry 211/212 General
Chemistry courses taught at George Mason are
intended for those students enrolled in a science
/engineering oriented curricula, with particular
emphasis on chemistry, biochemistry, and biology
The material on these slides is taken primarily
from the course text but the instructor has
modified, condensed, or otherwise reorganized
selected material.Additional material from other
sources may also be included. Interpretation of
course material to clarify concepts and solutions
to problems is the sole responsibility of this
instructor.
2
Thermochemistry

• Whenever matter changes composition, such as in a
  chemical reaction, the energy content of the
  matter changes also
• In some reactions the energy that the reactants
  contain is greater than the energy contained by
  the products
• This excess energy is released as heat
• In other reactions it is necessary to add energy
  (heat) before the reaction can proceed
• The energy contained by the products in these
  reactions is greater than the energy of the
  original reactants
• Physical changes can also involve a change in
  energy such as when ice melts

3
Thermochemistry

• Thermodynamics is the science of the relationship
  between heat and other forms of energy
• Thermochemistry is the study of the quantity of
  heat absorbed or evolved by chemical reactions
• Energy is the potential or capacity to move
  matter (do work) energy is a property of matter
• Energy can be in many forms
• Radiant Energy - Electromagnetic radiation
• Thermal Energy - Associated with random motion
  of a molecule or atom
• Chemical Energy - Energy stored within the
  structural limits of a molecule or atom

4
Energy

• There are three broad concepts of energy
• Kinetic Energy (Ek) is the energy associated with
  an object by virtue of its motion,
• Ek ½mv2
• Potential Energy (Ep) is the energy an object has
  by virtue of its position in a field of force,
• Ep mgh
• Internal Energy (Ei or Ui) is the sum of the
  kinetic and potential energies of the particles
  making up a substance
• E(i) Ek Ep

5
Energy

• SI unit of energy is the Joule (J) kg?m2/s2
• 1 watt 1 J/s
• 1 cal amount of energy needed to raise 1 g
  of water 1 oC (common energy unit)
• 1 cal 4.181 J
• 1 Btu 1055 J (Btu - British Thermal Unit)
• The Law of Conservation of Energy Energy may be
  converted from one form to another, but the total
  quantities of energy remain constant

6
Practice Problem

• Thermal decomposition of 5.0 metric tons of
  limestone (CaCO3) and Carbon Dioxide (CO2)
  requires 9.0 x 106 kJ of heat.
• Convert this energy to
• a. Joules b. calories c. British
  Thermal Units (Btu)

7
Energy

• When a chemical system changes from reactants to
  products and the products are allowed to return
  to the starting temperature, the Internal Energy
  (E) has changed (?E)
• The difference between the system internal energy
  after the change (Efinal) and before the change
  (Einitial) is
• ?E Efinal - Einitial Eproducts -
  Ereactants
• If energy is lost to the surroundings, then
• Efinal lt Einitial ??E lt 0
  
• If energy is gained from the surroundings, then
• Efinal gt Einitial ??E gt 0

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Heat of Reaction

• Heat of a system is denoted by the symbol q
• The sign of q is positive if heat is absorbed by
  the system, i.e., temperature increases
• The sign of q is negative if heat is evolved by
  the system, i.e., temperature decreases
• Heat of Reaction is the value of q required to
  return a system to the given temperature at the
  completion of the reaction

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Heat of Reaction

• An Exothermic (heat out) process releases
  (evolves) heat decreasing Enthalpy
• (q lt 0 negative)
• An Endothermic (heat in) process absorbs heat
  from the surroundings increasing the Enthalpy
• (q gt 0 positive)

Exothermic
Endothermic
qlt0
qgt0
Energy
Energy
Surroundings
Surroundings
System
System
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Kinetic-theory Explanation of Heat

• Heat motion on a molecular scale

Heat Energy
Higher velocity on impact
Lower velocity on impact
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Heat of Reaction

• In chemical reactions, heat is often transferred
  from the system to its surroundings, or vice
  versa
• The substance or mixture of substances under
  study in which a change occurs is called the
  thermodynamic system (or simply system)
• The surroundings are everything in the vicinity
  of the thermodynamic system
• Heat is defined as the energy that flows into or
  out of a system because of a difference in
  temperature between the system and its
  surroundings
• Heat flows from a region of higher temperature
  to one of lower temperature
• Once the temperatures become equal, heat flow
  stops

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Heat Flow and Phase Changes

• Predict the sign of q for each of the processes
  below
• 1. H2O(g) ? H2O(l)
• 2. CO2(s) ? CO2(g)
• 3. CH4(g) O2(g) ? CO2(g) H2O(g)

Condensation - Energy (heat) is lost by water
vapor ? q is negative
Evaporation Energy (heat) is absorbed (added)
by the system) ? q is positive
Combustion Burning (oxidation) of Methane is an
exothermic reaction heat
is evolved ? q is negative
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Work Internal Energy

• Internal Energy
• The Internal Energy of a system, E, is precisely
  defined as the heat at constant pressure (qp)
  plus the work (w) done by the system
• Work is the energy transferred when an object is
  moved by a force

Internal Energy used to expand volume by
increasing pressure is lost to the surroundings,
thus the negative sign

• Adiabatic Process Thermodynamic Process
  Without the Gain or Loss of Heat (?q 0)

14
Practice Problem

• A system delivers 225 J of heat to the
  surroundings while delivering 645 J of
  work.Calculate the change in the internal
  energy, ?E, of the system

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Pressure-Volume Work
Sign conventions for q, w, and ?E q
w (-P?V) ?E

Depends on sizes of q
and w
Depends on sizes of q and w

q system gains heat q system loses
heat w work done on system w work
done by system
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Practice Problem

• A system expands in volume from 2.0 L to 24.5 L
  at constant temperature.
• Calculate the work (w), in Joules (J), if the
  expansion occurs against a constant pressure of
  5.00 atm

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Practice Problem

• A system that does no work but which transfers
  heat to the surrounding has
• a. q lt 0, ?E gt 0 b. q lt 0, ?E lt 0
  c. q gt 0, ?E gt 0 d. q gt 0, ?E lt
  0 e. q lt 0, ?E 0
• A system that does no work but receives heat
  from the surroundings has
• a. q lt 0, ?E gt 0 b. q gt 0, ?E lt 0
  c. q ?E d. q -
  ?E e. w ?E
• A system which undergoes an adiabatic change
  (i.e., ?q 0) and does work on the surroundings
  has
• a. w lt 0, ?E 0, b. w gt 0, ?E gt 0
  c. w gt 0, ?E lt 0 d. w lt 0, ?E gt
  0 e. w lt 0, ?E lt 0
• A system which undergoes an adiabatic change
  (i.e., ?q 0) and has work done on it by the
  surroundings has
• a. w ?E b. w -?E
  c. w gt 0, ?E lt 0 d. w lt 0, ?E
  gt 0 e. w gt ?E

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Enthalpy and Enthalpy Change

• Enthalpy, denoted H, is an extensive property of
  a substance that can be used to obtain the heat
  absorbed or evolved in a chemical reaction
• An extensive property is one that depends on the
  quantity of substance
• Enthalpy is a state function, a property of a
  system that depends only on its present state and
  is independent of any previous history of the
  system
• Enthalpy represents the heat energy tied up in
  chemical bonds

19
Enthalpy and Enthalpy Change

• The change in Enthalpy for a reaction at a given
  temperature and pressure (called the Enthalpy of
  Reaction) is obtained by subtracting the enthalpy
  of the reactants from the enthalpy of the
  products.

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Enthalpy and Enthalpy Change

• Enthalpy is defined as the internal energy plus
  the product of the pressure and volume (work)
• The change in Enthalpy is the change in internal
  energy plus the product of constant pressure and
  the change in Volume

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Enthalpy and Enthalpy Change

• Recall
• The change in Enthalpy equals the heat gained or
  lost (heat of reaction) at constant pressure
• This represents the entire change in internal
  energy (DE) minus any expansion work done by
  the system (P?V would have negative sign)

(At Constant Pressure)
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Practice Problem

• An ideal gas (the system) is contained in a
  flexible balloon at a pressure of 1 atm and is
  initially at a temperature of 20.0oC.
• The surrounding air is at the same pressure, but
  its temperature is 25oC. When the system is
  equilibrated with its surroundings, both systems
  and surroundings are at 25oC and 1 atm.
• In changing from the initial to the final state,
  which of the following relationships regarding
  the system is correct?
• ?E 0
• ?E lt 0
• ?H 0
• w gt 0
• q gt 0

Heat is added, internal energy increases
? ?E gt 0
? ?E gt 0
Heat is added, internal energy increases
? ?H gt 0
?E increases and P?V work is done by system
? W lt 0
P?V work is done by system (volume increase)
Temperature (heat) in system increases
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Practice Problem

• In which of the following processes is ?H ?E?
• a. 2HI(g) ? H2(g) I2(g) at
  atmospheric pressure
• (P?V 0 no change in moles, volume)
• b. Two moles of ammonia gas are cooled from
  325oC to 300oC at 1.2 atm
• (P?V ? 0 Vol decreases)
• c. H2O(l) ? H2O(g) at 100oC at atmospheric
  pressure
• (P?V ? 0 Vol increases)
• d. CaCO3(s) ? CaO(s) CO2 (g) at 800oC
  at atmospheric pressure
• (P?V ? 0 Vol increases)
• e. CO2(s) ? CO2(g) at atmospheric pressure
• (P?V ? 0 Vol increases)

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Comparing ?E ?H

• Reactions that do not involve gases
• Reactions such as precipitation, acid-base, many
  redox, etc., do not produce gases
• Since the change in volumes of liquids and
  solids are quite small
• ?V ? 0 P ?V ? 0 ?H ? ?E
• Reactions in which the amount (mol) of gas does
  not change
• (Vol of Gaseous Reactants Vol Gaseous Products
• ?V 0 P ?V 0 ?H ?E
• Reactions in which the amount (mol) of gas does
  change
• P?V ? 0
• However, qp is usually much greater than
  P?V
• Therefore ?H ? ?E

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Comparing ?E ?H

• Example
• 2H2(g) O2(g) ? 2H2O(g)
• Change in moles 3 mol ? 2 mol ? P?V ? 0
• ?H -483.6 kJ and P?V -2.5kJ
• ?E ?H - P?V -483.6 kJ - (-2.5 kJ)
  -481.1 kJ
• Most of ?E occurs as Heat (?H qp)
• ?H ? ?E
• For many reactions, even when P?V ? 0, ?H is
  close to ?E

26
Comparing ?E ?H

• For which one of the following reactions will ?H
  be approximately (or exactly) equal to ?E?
• a. H2(g) Br2(g) ? 2HBr(g)
• (No change in volume no change in work, P?V
  0)
• b. H2 O(l) ? H2O(g)
• (Change in volume change in work due to gas
  expansion, P?V ? 0)
• c. CaCO3(s) ? CaO(s) CO2(g)
• (Change in volume change in work due to gas
  expansion, P?V ? 0
• d. 2H(g) O(g) ? H2O(l)
• (Change in volume condensation, heat (q)
  released, P?V ? 0)
• e. CH4(g) 2O2(g) ? CO2(g)
  2H2O(l)
• (Change in volume condensation, heat (q)
  released, P?V ? 0)

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Exothermic and Endothermic Processes

• Energy (E), Pressure (P), and Volume (V) are
  state functions
• Enthalpy (H) is also a state function, which
  means that ?H depends only on the difference
  between Hfinal Hinitial
• The Enthalpy change of a reaction, also called
  the Heat of Reaction (?Hrxn), always refers to
• ?Hrxn ?Hfinal - ?Hinitial
  ?Hproducts - ?Hreactants
• ?Hproducts can be either more or less than
  ?Hreactants
• The resulting sign of ?H indicates whether heat
  is absorbed from the surroundings (heat in) or
  released to the surroundings (heat out) in the
  process

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Exothermic and Endothermic Processes

• An Exothermic reaction releases heat (heat out)
  to surroundings with a decrease in system
  Enthalpy
• CH4(g) 2O2 ? CO2(g) 2H2O(g)
  heat
• Exothermic ?Hfinal lt ?Hinitial
  ?H lt 0 (negative)
• An Endothermic reaction absorbs heat (heat in)
  from the surroundings resulting in an increase in
  system Enthalpy
• Heat H2O(s) ? H2O(l)
• Endothermic ?Hfinal gt ?Hinitial
  ?H gt 0 (positive)

29
Types of Enthalpy Changes

• When a compound is produced from its elements,
  the Enthalpy change (Heat of Reaction) is called
• Heat of Formation (?Hf)
• K(s) ½Br2()l) ? KBr(s) ?H ?Hf
• When a substance melts, the Enthalpy change is
  called
• Heat of Fusion (?Hfus)
• NaCl(s) ? NaCl(l) ?H ?H(fus)
• When a substance vaporizes, the Enthalpy change
  is called
• Heat of Vaporization
• C6H6(l) ? C6H6(g) ?H ?H(vap)

30
Thermochemical Equations

• A thermochemical equation is the chemical
  equation for a reaction (including phase labels)
  in which the equation is given a molar
  interpretation, and the Enthalpy of Reaction for
  these molar amounts is written directly after the
  equation.

?H is negative heat is lost to surroundings 1
mol N2 3 mol H2 yields 91.8 kJ of heat
31
Practice Problem

• Sulfur, S8, burns in air to produce sulfur
  dioxide. The reaction evolves (releases) 9.31 kJ
  of heat per gram of sulfur at constant pressure.
  Write the thermochemical equation for this
  reaction.

32
Practice Problem

• In a phase change of water between the liquid and
  the gas phases, 770.1 kJ of energy was released
  by the system. What was the product, and how
  much of it was formed in the phase change.
• (Data H2O(l) ? H2O(g) ?H 44.01
  kJ/mol)
• a. 315 g of water vapor was produced
• b. 17.5 g of water vapor was produced
• c. 17.5 mol of water vapor was produced
• d. 17.5 mol of liquid water was produced
• ?H is positive (endothermic reaction)Since
  energy was released, the gas condensed to liquid
  
• e. 17.5 g of liquid water was produced

33
Thermochemical Equations

• The following are two important rules for
  manipulating Thermochemical equations
• When a thermochemical equation is multiplied by
  any factor, the value of ?H for the new equation
  is obtained by multiplying the ?H in the original
  equation by that same factor
• When a chemical equation is reversed, the value
  of ?H is reversed in sign

34
Practice Problem

• When white Phosphorus burns in air, it produces
  Phosphorus (V) Oxide (Change in Oxidation state)
• P4(s) 5O2(g) ? P4O10(s) ?H -3010 kJ
• What is ?H for the following equation?
• P4O10(s) ? P4(s) 5O2(g) ?H ?
• Ans
• The original reaction is reversed
• ? change the sign
• ?H 3010 kJ

35
Practice Problem

• Carbon Disulfide (CS2(l)) burns in air, producing
  Carbon Dioxide and Sulfur Dioxide
• CS2(l) 3O2(g) ? CO2(g) 2SO2(g) ?H
  -1077 kJ
• What is ?H for the following equation?
• 1/2 CS2(l) 3/2 O2(g) ? 1/2 CO2(g)
  SO2(g)
• Ans The reaction uses ½ the amounts
• ? Divide ?H by 2
• ?H (-1077 / 2) - 538.5 kJ

36
Applying Stoichiometry andHeats of Reactions

• Consider the reaction of Methane, CH4, burning in
  the presence of Oxygen at constant pressure.
  Given the following equation, how much heat could
  be obtained by the combustion of10.0 grams CH4?

37
Measuring Heats of Reaction

• To see how Heats of Reactions (Enthalpy change of
  reaction ?H) are measured, we must look at the
  heat required to raise the temperature of a
  substance
• A thermochemical measurement is based on the
  relationship between heat and temperature change
• The heat required to raise the temperature of a
  substance is its Heat Capacity

38
Measuring Heats of Reaction

• Heat Capacity and Specific Heat
• The molar heat capacity, C, of a sample of
  substance is the quantity of heat required to
  raise the temperature of one mole of substance
  one degree Celsius
• C is in units of J/mol ? oC, n moles of
  substance
• The specific heat capacity, S, (or specific
  heat) is the heat required to raise the
  temperature ofone gram of a substance by one
  degree Celsius
• S is in units of J/g ? oC m grams
  of sample

?T Tfinal - Tinitial
?T Tfinal - Tinitial
39
Measuring Heats of Reaction
40
Practice Problem
Suppose you mix 20.5 g of water at 66.2 oC with
45.4 g of water at 35.7 oC in an insulated cup.
What is the maximum temperature of the solution
after mixing? Ans The heat lost by the water at
66.2 oC is balanced by the heat gained by the
water at 35.7 oC
41
Measuring Heats of Reaction

• Bomb Calorimeter used to measure heats of
  combustion

42
Practice Problem

• How much heat is gained by Nickel when 500 g of
  Nickel is warmed from 22.4 to 58.4C?
• The specific heat of Nickel is 0.444 J/(g C)
• a. 2000 J b. 4000 J c. 6000 J
• d. 8000 J e. 10000 J
• Ans d

43
Practice Problem

• When 25.0 mL of 0.5 M H2SO4 is added to 25.0 mL
  of 1.00 M KOH in a calorimeter at 23.5 oC, the
  temperature rises to 30.17oC
• Calculate ?Hrxn for each reactant. Assume density
  (d) and specific heatof the solution (s) are the
  same as water

Cont on next Slide
44
Practice Problem (Cont)

• When 25.0 mL of 0.5 M H2SO4 is added to 25.0 mL
  of 1.00 M KOH in a calorimeter at 23.5 oC, the
  temperature rises to 30.17oC.
• Calculate ?Hrxn for each reactant. Assume density
  (d) and specific heat of the solution are the
  same as water.

Temperature of water increased (23.5oC ?
30.17oC) ?The Reaction is Exothermic (heat
released to surroundings (water)) Thus, qrxn is
negative
45
Hesss Law

• Hesss law of Heat Summation
• For a chemical equation that can be writtenas
  the sum of two or more steps, theEnthalpy change
  for the overall equationis the sum of the
  Enthalpy changes for the individual steps
• In coupled reactions, the Enthalpy change for the
  overall reaction is the sum of the Enthalpy
  changes for the coupled reactions
• Often need to reverse chemical equations to
  couple them so chemical species are on the
  correct side of yield sign, or multiply through
  by a coefficient to cancel common chemical
  species

46
Hesss Law

• For example, suppose you are given the following
  data

Could you use these data to obtain the enthalpy
change for the following reaction?
Cont on next Slide
47
Hesss Law

• If we multiply the first equation by 2 and
  reverse the second equation, they will sum
  together to become the third

Note the change in ?H values with the changes in
the molar coefficients to balance the first
equation and the reversal of equation 2
48
Practice Problem

• Given the following data,
• A(s) O2(g) ? AO2(g) ?H 105 kJ/mol
• A(g) O2(g) ? AO2(g) ?H 1200 kJ/mol
• Find the heat required for the reaction
  converting
• A(s) to A(g) at 298 K and 1 atm
  pressure.

49
Standard Enthalpies of Formation

• The term standard state refers to the standard
  thermodynamic conditions chosen for substances
  when listing or comparing thermodynamic data
• Pressure - 1 atmosphere (760 mm Hg)
• Temperature - (usually 25oC).
• The Enthalpy change for a reaction in which
  reactants are in their standard states is denoted
  as the Standard Heat of Reaction

50
Standard Enthalpies of Formation

• Standard Enthalpy of Formation of Substance
• The Enthalpy change for the formation ofone mole
  of a substance in its standard state from its
  component elements in their standard states
• Note The standard Enthalpy of Formation for a
• Pure Element (C, Fe, Au, N)
• in its standard state is zero

51
Standard Enthalpies of Formation

• Law of Summation of Heats of Formation
• The Enthalpy of a reaction i.e., the Standard
  Heat of Reaction
• (?Horxn)
• is equal to the total formation energy of the
  products minus that of the reactants
• Where ? is the mathematical symbol meaning
• the sum of
• and m and n are the coefficients of the
  substances in the chemical equation, i.e., the
  relative number of moles of each substance

52
Standard Enthalpies of Formation
Selected Standard Heats of Formation
(Enthalpies) At 25oC (298oK)
53
Practice Problem

• Calculate the Heat of Reaction, ?Hrxn, for the
  combustion of C3H6(g)
• C3H6(g) 9/2 O2(g) ? 3 CO2(g) 3 H2O(l)
• ?Hof values in kilojoules per mole are as
  follows
• C3H6(g) 21 CO2(g) 394 H2O(l)
  286
• a. 2061 kJ b. 2019 kJ c. 701 kJ
• d. 2019 kJ e. 2061 kJ
• Ans a

54
Practice Problem

• Acetylene burns in air according to the equation
  below.
• Given ?Hof CO2(g) -395.5 kJ/mol
• ?Hof H2O(g) -241.8 kJ/mol

Calculate ?Hof of C2H2(g)
55
Summary Equations Relationships