Chapter 5 discusses thermochemistry, focusing on kinetic and potential energy, heat capacity, and the laws governing energy transfer in chemical reactions. It introduces key concepts such as enthalpy, calorimetry, and the principles of the first law of thermodynamics, emphasizing energy conservation and transformation. The document explains how changes in energy are related to heat and work in different types of systems.

1. Chapter 5
Thermochemistry

2. 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


3. 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

4. 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

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

6. Constant Pressure
Calorimetry

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

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

9. 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

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

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

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

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

14. 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

15. Enthalpies of formation

16. • 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 

18. • 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



19. • 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 


20. Enthalpy

21. • 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



22. • 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

23. 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 

24. 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 

25. • 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 

27. Sign Conventions