Chapter 5
Thermochemistry

Kinetic Energy and Potential Energy
• Energy is usually defi ned as the capacity to do work
• Kinetic energy is the energy of motion:
• Potential energy is the energy an object possesses by virtue of its position.
• Potential energy can be converted into kinetic energy. Example: a bicyclist at the
top of a hill.
• Electrostatic potential energy, Ed, is the attraction between two oppositely
charged particles, Q1 and Q2, a distance d apart:
• The constant  = 8.99  109 J-m/C2.
• If the two particles are of opposite charge, then Ed is the electrostatic attraction
between them.
2
2
1
mvEk 
The Nature of Energy
d
QQ
Ed
21


Units of Energy
• SI Unit for energy is the joule, J:
• We sometimes use the calorie instead of the joule:
1 cal = 4.184 J (exactly)
• A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal
  
J1m/s-kg1
m/s1kg2
2
1
2
1
2
22

 mvEk

Heat Capacity and Specific Heat
• Calorimetry = measurement of heat flow.
• Calorimeter = apparatus that measures heat flow.
• Heat capacity = the amount of energy required to raise
the temperature of an object (by one degree).
• Molar heat capacity = heat capacity of 1 mol of a
substance.
• Specific heat = specific heat capacity = heat capacity of 1
g of a substance.
   
if TTT
Tq

 substanceofgramsheatspecific
Calorimetry

• Atmospheric pressure is constant!
Constant Pressure Calorimetry
 
  T
qq


solutionofgrams
solutionofheatspecificsolnrxn

Constant Pressure
Calorimetry

• Reaction carried out
under constant volume.
• Use a bomb
calorimeter.
• Usually study
combustion.
Bomb Calorimetry (Constant Volume
Calorimetry)
TCq  calrxn
Example 5.1

Systems and Surroundings
• Analyzing Energy Changes
System: part of the universe we are interested in.
Surroundings: the rest of the universe.

Types of systems
A. open: exchange of matter and
energy
- example: living systems
E0, m0
B. closed: exchange of energy, no
exchange of matter
– Example: photochemical
batch reactor.
E0, m=0
C. isolated: no interaction at all with
surroundings- - Example:
Dewar flask with absolutely rigid
walls in a Faraday cage on a
space station
idealization, but useful
E=0, m=0

Thermochemical Equations
• Enthalpy, H: Heat transferred between the system
and surroundings carried out under constant
pressure.
• Enthalpies of reactions
• H = Hproducts - Hreactants

Exothermic and Endothermic Processes
a. Endothermic:
absorbs heat from
the surroundings.
• An endothermic
reaction feels cold.

Thermochemical Equations
b. Exothermic:
transfers heat to
the surroundings.
• An exothermic
reaction feels hot.

State Functions
• State function: depends only on the initial and final
states of system, not on how the internal energy is
used.

The Law of Hess
• When reactants are converted to
products, the change in enthalpy is
the same whether the reaction
takes place in one step or in a
series of steps.
• Hess’s law is based on the fact
that because H is a state function,
H depends only on the initial
and final state (that is, only on the
nature of reactants and products).
• The enthalpy change would be the
same whether the overall reaction
takes place in one step or many
steps.
• Example 5.3, page 98

Enthalpies of formation

• Chemical reactions can absorb or release heat.
• However, they also have the ability to do work.
• For example, when a gas is produced, then the gas
produced can be used to push a piston, thus doing work.
Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)
• The work performed by the above reaction is called
pressure-volume work.
• When the pressure is constant,
Work
VPw 

• Enthalpy, H: Heat transferred between the system and
surroundings carried out under constant pressure.
• Enthalpy is a state function.
• If the process occurs at constant pressure,
Enthalpy
PVEH 
 
VPE
PVEH



• Since we know that
• We can write
• When H, is positive, the system gains heat from the
surroundings.
• When H, is negative, the surroundings gain heat from
the system.
Enthalpy
VPw 
wq
VPEH
P 


Enthalpy

• For a reaction:
• Enthalpy is an extensive property (magnitude H is
directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)DH = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = -1604 kJ
Enthalpies of Reaction
reactantsproducts
initialfinal
HH
HHH



• When we reverse a reaction, we change the sign of DH:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g)DH = +802 kJ
• Change in enthalpy depends on state:
H2O(g)  H2O(l) DH = -88 kJ
Enthalpies of Reaction

Transferring Energy
Work and Heat
• Force is a push or pull on an object.
• Work is the product of force applied to an object over a distance:
• Energy is the work done to move an object against a force.
• Heat is the transfer of energy between two objects.
• Energy is the capacity to do work or transfer heat.
dFw 

The First Law of Thermodynamics
Energy is neither created or destroyed
• Total energy lost by a system equals the total energy gained by a system.
• Internal Energy: total energy of a system (kinetic + potential).
• Cannot measure absolute internal energy.
• Change in internal energy,
initialfinal EEE 

• Energy of (system + surroundings) is constant.
• Any energy transferred from a system must be
transferred to the surroundings (and vice versa).
• From the first law of thermodynamics:
when a system undergoes a physical or chemical change, the
change in internal energy is given by the heat added to or
absorbed by the system plus the work done on or by the system:
Relating E to Heat and Work
wqE 

Sign Conventions