Download Chemistry: McMurry and Fay, 5th Edition

John E. McMurry • Robert C. Fay
C H E M I S T R Y
Fifth Edition
Chapter 8
Thermochemistry: Chemical Energy
Lecture Notes
Alan D. Earhart
Southeast Community College • Lincoln, NE
Copyright © 2008 Pearson Prentice Hall, Inc.
Energy
Energy: The capacity to supply heat or do work.
Kinetic Energy (EK): The energy of motion.
Potential Energy (EP): Stored energy.
Units:
1 cal = 4.184 J (exactly)
1 Cal = 1000 cal = 1 kcal
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Chapter 8/2
Energy Changes and Energy
Conservation
Law of Conservation of Energy: Energy cannot be
created or destroyed; it can only be converted from one
form to another.
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Chapter 8/3
Energy Changes and Energy
Conservation
Total Energy = EK + EP
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Chapter 8/4
Energy Changes and Energy
Conservation
Thermal Energy: The kinetic energy of molecular
motion and is measured by finding the temperature of
an object.
Heat: The amount of thermal energy transferred from
one object to another as the result of a temperature
difference between the two.
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Chapter 8/5
Internal Energy and State
Functions
First Law of Thermodynamics: The total internal
energy of an isolated system is constant.
DE = Efinal - Einitial
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Chapter 8/6
Internal Energy and State
Functions
First Law of Thermodynamics: The total internal
energy of an isolated system is constant.
DE = Efinal - Einitial
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Chapter 8/7
Internal Energy and State
Functions
First Law of Thermodynamics: The total internal
energy of an isolated system is constant.
DE = Efinal - Einitial
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Chapter 8/8
Internal Energy and State
Functions
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g) + 802 kJ energy
DE = Efinal - Einitial = -802 kJ
802 kJ is released when 1 mole of methane, CH4, reacts
with 2 moles of oxygen to produce 1 mole of carbon
dioxide and two moles of water.
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Chapter 8/9
Internal Energy and State
Functions
State Function: A function or property whose value
depends only on the present state, or condition, of the
system, not on the path used to arrive at that state.
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Chapter 8/10
Expansion Work
Work = Force x Distance
w=Fxd
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Chapter 8/11
Expansion Work
C3H8(g) + 5O2(g)
3CO2(g) + 4H2O(g)
6 mol of gas
7 mol of gas
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Chapter 8/12
Expansion Work
Expansion Work: Work done as the result of a volume
change in the system. Also called pressure-volume or
PV work.
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Chapter 8/13
Energy and Enthalpy
DE = q + w
q = heat transferred
w = work = -PDV
q = DE + PDV
Constant Volume (DV = 0): qV = DE
Constant Pressure: qP = DE + PDV
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Chapter 8/14
Energy and Enthalpy
qP = DE + PDV = DH
Enthalpy change
or
Heat of reaction (at constant pressure)
Enthalpy is a state function whose
value depends only on the current
state of the system, not on the
path taken to arrive at that state.
DH = Hfinal - Hinitial
= Hproducts - Hreactants
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Chapter 8/15
The Thermodynamic Standard
State
C3H8(g) + 5O2(g)
3CO2(g) + 4H2O(g)
DH = -2043 kJ
C3H8(g) + 5O2(g)
3CO2(g) + 4H2O(l)
DH = -2219 kJ
Thermodynamic Standard State: Most stable form of a
substance at 1 atm pressure and at a specified
temperature, usually 25 °C; 1 M concentration for all
substances in solution.
C3H8(g) + 5O2(g)
3CO2(g) + 4H2O(g)
DH° = -2043 kJ
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Chapter 8/16
Enthalpies of Physical and
Chemical Change
Enthalpy of Fusion (DHfusion): The amount of heat
necessary to melt a substance without changing its
temperature.
Enthalpy of Vaporization (DHvap): The amount of heat
required to vaporize a substance without changing its
temperature.
Enthalpy of Sublimation (DHsubl): The amount of heat
required to convert a substance from a solid to a gas
without going through a liquid phase.
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Chapter 8/17
Enthalpies of Physical and
Chemical Change
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Chapter 8/18
Enthalpies of Physical and
Chemical Change
2Al(s) + Fe2O3(s)
2Fe(s) + Al2O3(s)
DH° = -852 kJ
exothermic
2Fe(s) + Al2O3(s)
2Al(s) + Fe2O3(s)
DH° = +852 kJ
endothermic
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Chapter 8/19
Calorimetry and Heat Capacity
Measure the heat flow at constant pressure (DH).
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Chapter 8/20
Calorimetry and Heat Capacity
Measure the heat flow at constant volume (DE).
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Chapter 8/21
Calorimetry and Heat Capacity
Heat Capacity (C): The amount of heat required to raise
the temperature of an object or substance a given
amount.
q = C x DT
Molar Heat Capacity (Cm): The amount of heat required
to raise the temperature of 1 mol of a substance by 1 °C.
q = (Cm) x (moles of substance) x DT
Specific Heat: The amount of heat required to raise the
temperature of 1 g of a substance by 1 °C.
q = (specific heat) x (mass of substance) x DT
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Chapter 8/22
Calorimetry and Heat Capacity
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Chapter 8/23
Calorimetry and Heat Capacity
Assuming that a can of soda has the same specific
heat as water, calculate the amount of heat (in
kilojoules) transferred when one can (about 350 g) is
cooled from 25 °C to 3 °C.
q = (specific heat) x (mass of substance) x DT
J
Specific heat = 4.18
g °C
Mass = 350 g
Temperature change = 25 °C - 3 °C = 22 °C
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Chapter 8/24
Calorimetry and Heat Capacity
Calculate the amount of heat transferred.
4.18 J x 350 g x 22 °C
Heat evolved =
= 32 000 J
g °C
32 000 J x
1 kJ
= 32 kJ
1000 J
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Chapter 8/25
Hess’s Law
Hess’s Law: The overall enthalpy change for a reaction
is equal to the sum of the enthalpy changes for the
individual steps in the reaction.
Haber Process: 3H2(g) + N2(g)
2NH3(g) DH° = -92.2 kJ
Multiple-Step Process
2H2(g) + N2(g)
N2H4(g)
DH°1 = ?
N2H4(g) + H2(g)
2NH3(g)
DH°2 = -187.6 kJ
3H2(g) + N2(g)
2NH3(g)
DH°1+2 = -92.2 kJ
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Chapter 8/26
Hess’s Law
DH°1 + DH°2 = DH°1+2
DH°1 = DH°1+2 - DH°2
= -92.2 kJ - (-187.6 kJ) = 95.4 kJ
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Chapter 8/27
Standard Heats of Formation
Standard Heat of Formation (DH°f ): The enthalpy
change for the formation of 1 mol of a substance in its
standard state from its constituent elements in their
standard states.
Standard states
C(s) + 2H2(g)
CH4(g)
DH°f = -74.8 kJ
1 mol of 1 substance
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Chapter 8/28
Standard Heats of Formation
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Chapter 8/29
Standard Heats of Formation
DH° = DH°f (Products) - DH°f (Reactants)
aA + bB
cC + dD
DH° = [c DH°f (C) + d DH°f (D)] - [a DH°f (A) + b DH°f (B)]
Products
Reactants
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Chapter 8/30
Standard Heats of Formation
Using standard heats of formation, calculate the
standard enthalpy of reaction for the photosynthesis of
glucose (C6H12O6) and O2 from CO2 and liquid H2O.
6CO2(g) + 6H2O(l)
C6H12O6(s) + 6O2(g)
DH° = ?
DH° = [DH°f (C6H12O6(s))] - [6 DH°f (CO2(g)) + 6 DH°f (H2O(l))]
DH° = [(1 mol)(-1260 kJ/mol)] [(6 mol)(-393.5 kJ/mol) + (6 mol)(-285.8 kJ/mol)]
= 2816 kJ
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Chapter 8/31
Standard Heats of Formation
6CO2(g) + 6H2O(l)
C6H12O6(s) + 6O2(g) DH° = 2816 kJ
Why does the calculation “work”?
1. Reverse the “reaction” and reverse the sign on
the standard enthalpy change.
C(g) + O2(g)
CO2(g)
DH° = -393.5 kJ
C(g) + O2(g)
DH° = 393.5 kJ
Becomes
CO2(g)
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Chapter 8/32
Standard Heats of Formation
6CO2(g) + 6H2O(l)
C6H12O6(s) + 6O2(g) DH° = 2816 kJ
Why does the calculation “work”?
2. Multiply the coefficients by a factor and multiply
the standard enthalpy change by the same factor.
CO2(g)
C(g) + O2(g)
DH° = -393.5 kJ
Becomes
6CO2(g)
6C(g) + 6O2(g) DH° = 6(393.5 kJ) = 2361 kJ
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Chapter 8/33
Standard Heats of Formation
6CO2(g) + 6H2O(l)
C6H12O6(s) + 6O2(g) DH° = 2816 kJ
Why does the calculation “work”?
C6H12O6(s)
DH° = -1260 kJ
6CO2(g)
6C(g) + 6O2(g)
DH° = 2361 kJ
6H2O(l)
6H2(g) + 3O2(g)
DH° = 1715 kJ
6C(s) + 6H2(g) + 3O2(g)
6CO2(g) + 6H2O(l)
C6H12O6(s) + 6O2(g) DH° = 2816 kJ
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Chapter 8/34
Bond Dissociation Energies
Bond dissociation energies are standard enthalpy
changes for the corresponding bond-breaking reactions.
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Chapter 8/35
Bond Dissociation Energies
H2(g) + Cl2(g)
2HCl(g)
DH° = D(Reactant bonds) - D(Product bonds)
DH° = (DH-H + DCl-Cl) - (2DH-Cl)
DH° = [(1 mol)(436 kJ/mol) + (1mol)(243 kJ/mol)] -
(2mol)(432 kJ/mol)
= -185 kJ
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Chapter 8/36
Fossil Fuels, Fuel Efficiency,
and Heats of Combustion
CH4(g) + 2O2(g)
CO2(g) + 2H2O(l)
DH°c = [DH°f (CO2(g)) + 2 DH°f (H2O(l))] - [DH°f (CH4(g))]
= [(1 mol)(-393.5 kJ/mol) + (2 mol)(-285.8 kJ/mol)] [(1 mol)(-74.8 kJ/mol)]
= -890.3 kJ
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Chapter 8/37
Fossil Fuels, Fuel Efficiency,
and Heats of Combustion
CH4(g) + 2O2(g)
CO2(g) + 2H2O(l)
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Chapter 8/38
An Introduction to Entropy
Spontaneous Process: A process that, once
started, proceeds on its own without a continuous
external influence.
Entropy (S): The amount of molecular
randomness in a system.
Spontaneous processes are
• favored by a decrease in H (negative DH).
• favored by an increase in S (positive DS).
Nonspontaneous processes are
• favored by an increase in H (positive DH).
• favored by a decrease in S (negative DS).
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Chapter 8/39
An Introduction to Free Energy
Gibbs Free Energy Change (DG)
DG = DH - T DS
Enthalpy of
reaction
Temperature
(Kelvin)
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Entropy
change
Chapter 8/40
An Introduction to Free Energy
Gibbs Free Energy Change (DG)
DG = DH - T DS
DG < 0
Process is spontaneous
DG = 0
Process is at equilibrium
(neither spontaneous nor nonspontaneous)
DG > 0
Process is nonspontaneous
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Chapter 8/41