Transcript Chapter 8 Thermochemistry: Chemical Energy

C H E M I S T R Y
Chapter 8
Thermochemistry: Chemical Energy
Energy and Its Conservation
Conservation of Energy Law: Energy cannot be created or destroyed; it can
only be converted from one form to another.
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
Energy and Its Conservation
Energy and Its 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
Internal Energy and State Functions
First Law of Thermodynamics: The total internal energy E of an
isolated system is constant
DE = Efinal - Einitial
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.
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
Expansion Work
C3H8(g) + 5O2(g)
6 mol of gas
w=Fxd
3CO2(g) + 4H2O(g)
7 mol of gas
Expansion Work
Expansion Work: Work done as the result of a volume change in the system
8.5 Energy and Enthalpy
DE = q + w
q = heat transferred
w = work = -PDV
q = D E + PD V
Constant Volume (DV = 0):
Constant Pressure:
qV = DE
qP = DE + PDV
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
8.6The 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.
Standard enthalpy of reaction is indicated by the symbol ΔHo
C3H8(g) + 5O2(g)
3CO2(g) + 4H2O(g)
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DH° = -2043 kJ
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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
Enthalpies of Physical and Chemical
Change
Enthalpies of Physical and Chemical
Change
2Al(s) + Fe2O3(s)
2Fe(s) + Al2O3(s)
DHo = -852 kJ
exothermic
2Fe(s) + Al2O3(s)
2Al(s) + Fe2O3(s)
DHo = +852 kJ
endothermic
Examples
 Identify each of the following processes as
endothermic or exothermic and indicate the sign of
ΔHo
 Sweat evaporating from your skin
 Water freezing in a freezer
 Wood burning in fire
Example
 An LP gas tank in a home barbeque contains 13.2 kg of
propane, C3H8. Calculate the heat (in kJ) associated
with the complete combustion of all of the propane in
the tank
C3H8(g) + CO2(g)  3 CO2(g) + 4 H2O(g)
ΔHo = -2044kJ
Example
 How much heat (in kJ) is evolved when 5.00g of
aluminum reacts with a stoichiometric amount of
Fe2O3?
2 Al(s) + Fe2O3(s)  2 Fe(s) + Al2O3(s) ΔHo = -852 kJ
Calorimetry and Heat Capacity
Measure the heat flow at constant pressure (DH).
Calorimetry and Heat Capacity
Measure the heat flow at constant volume (DE).
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
Calorimetry and Heat Capacity
Molar Heat Capacity (Cm): The amount of heat necessary to raise the temperature
of 1 mol of a substance by 1 oC
q = Cm x Moles of substance x DT
Calorimetry and Heat Capacity
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Examples
 Assuming that Coca Cola has the same specific heat as
water [4.183 J/goC], calculate the amount of heat (in kJ)
transferred when one can (about 350.0 g) is cooled
from 25.0oC – 3.0oC
8.9 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°total = ???
Multiple-Step Process - Given
2H2(g) + N2(g)
N2H4(g)
N2H4(g) + H2(g)
2NH3(g)
3H2(g) + N2(g)
2NH3(g)
DH°1 = 95.4 Kj
DH°2 = -187.6 kJ
DH°total = DH°1 + DH°2
Hess’s Law
Example
 Find ΔHorxn for the following reaction:
3H2(g) + O3(g)  3 H2O(g)
ΔHorxn = ??
Use the following reactions with known ΔH’s
2H2 (g) + O2(g)  2 H2O(g)
Δ Ho = -483.6 kJ
3O2(g)  2 O3 (g)
Δ Ho = +285.4 kJ
Example
 Fin ΔHorxn for the following reaction
C(s) + H2O(g)  CO(g) + H2(g)
Horxn = ?
Use the following reactions with known H’s
C(s) + O2(g)  CO2(g)
ΔHo = -393.5 kJ
2CO(g) + O2(g)  2CO2(g)
Δ Ho = -566.0kJ
2H2 (g) + O2(g)  2H2O (g)
Δ Ho = -483.6 kJ
Standard Heats of Formation
Standard Heat of Formation (DHof ): 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)
1 mol of 1 substance
DHof = -74.8 kJ
Standard Heats of Formation
Standard Heats of Formation
DHo = DHof (Products) - DHof (Reactants)
aA + bB
cC + dD
DHo = [c DHof (C) + d DHof (D)] - [a DHof (A) + b DHof (B)]
Products
Reactants
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)
DHo = ?
Example
 Use the information in Table 8.2 to calculate ΔHo (in
kJ) for the reaction of ammonia with oxygen gas to
yield nitric oxide (NO) and water vapor, a step in the
Ostwald process for the commercial production of
nitric acid
An Introduction to Entropy
Spontaneous Process: A process that, once started, proceeds on its own
without a continuous external influence
An Introduction to Entropy
Entropy (S): The amount of molecular randomness in a system
An Introduction to Entropy
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).
Example
 Predict whether ΔSo is likely to be positive or negative
for each of the following reactions
 H2C=CH2(g) + Br2(g)  BrCH2CH2Br(l)
8.14 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
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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
Example
 Is the Haber process for the industrial synthesis of
ammonia spontaneous or nonspontaneous under
standard conditions at 25.0oC. At what temperature
(oC) does the changeover occur?
An Introduction to Free Energy
Gibbs Free Energy Change (DG)
DG = DH - T DS
Enthalpy of
reaction
Temperature
(Kelvin)
Entropy
change
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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