Chapter 8
Thermochemistry: Chemical
Energy
8.1 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 and Its Conservation
Energy: The capacity to supply heat or do work.
Kinetic Energy (EK): The energy of motion.
Potential Energy (EP): Stored energy.
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
Units:
1 cal = 4.184 J (exactly)
1 Cal = 1000 cal = 1 kcal
8.2 Internal Energy and State Functions
First Law of Thermodynamics: The total internal energy E of an
isolated system is constant
- The sum of all the kinetic and potential energies of particles making
up a substance
DE = Efinal - Einitial
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
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.
8.3 Expansion Work
Expansion Work: Work done as the result of a volume change
in the system
C3H8(g) + 5O2(g)
6 mol of gas
3CO2(g) + 4H2O(g)
7 mol of gas
w=Fxd
= - P∆V
8.4 Energy and Enthalpy
DE = q + w
q = heat transferred
q = DE + PDV
Constant Volume (DV = 0):
Constant Pressure:
w = work = -PDV
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)
DH° = -2043 kJ
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)
exothermic
The reverse reaction
2Fe(s) + Al2O3(s)
DHo = -852 kJ
2Al(s) + Fe2O3(s)
DHo = +852 kJ
endothermic
Applying Stoichiometry to Heats of
Reaction

a.
b.
c.
A propellant for rockets is obtained by mixing the liquids hydrazine,
N2H4, and dinitrogen tetroxide, N2O4. These compounds react to
give gaseous nitrogen, N2 and water vapor, evolving 1049 kJ of heat at
constant pressure when 1 mol N2O4 reacts.
Write the thermochemical equation for this reaction
Write the thermochemical equation for the reverse of the reaction
How much heat evolves when 10.0 g of hydrazine reacts according
to the reaction described in (a) ?
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
8.6Calorimetry 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
Example

Iron metal has a specific heat of 0.449 J/g oC. How much heat (in kJ)
is transferred to a 5.00g piece of iron, initally at 20.0oC, when it is
placed in a pot of boiling water? Assume that the temperature of the
water is 100.0oC and that the water remains at this temperature, which
is the final temperature of iron
Example

When 1.045 g of CaO is added to 50.0 mL of water at 25.0oC in a
calorimeter, the temperature of the water increases to 32.2 oC.
Assuming that the specific heat of the solution is 4.18 J/goC and that
the calorimeter itself absorbes a negligible amount of heat, calculate
∆H in kilojoules for the reaction
CaO(s) + H2O(l)  Ca(OH)2(aq)
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)
3O2(g)  2 O3 (g)
Δ Ho = -483.6 kJ
Δ Ho = +285.4 kJ
Example

Find 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)
2CO(g) + O2(g)  2CO2(g)
2H2 (g) + O2(g)  2H2O (g)
ΔHo = -393.5 kJ
Δ Ho = -566.0kJ
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

Calculate the enthalpy change for the following reaction
3 NO2(g) + H2O(l)  2 HNO3(aq) + NO(g)
*Use standard enthalpies of formation
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
a.
b.
H2C=CH2(g) + Br2(g)  BrCH2CH2Br(l)
Consider the following figures
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?
3H2(g) + N2(g)
2NH3(g)
DH°total = ???