1. Kinetics and Equilibrium
Unit VI

2. I Kinetics
A. Kinetics is the study of the rates of reactions and reaction mechanisms
Rate
Speed of a reaction
Mechanism
Steps involved in a reaction

3. B. Role of Energy in Kinetics
Activation Energy (Ea)
Amount of energy needed to start up a reaction
Heat of Reaction (H)
Measures the difference between the energy of the products and reactants in a reaction
Hreaction = Hproducts –Hreactants
Reference Table I
HNO2 = kJ
HNH3 = kJ

4. Reactions that release energy
C. Types of Reactions
Exothermic
Reactions that release energy
Energy is a product of the reaction
N2 + 3H2→ 2NH kJ
 H is negative ( H= kJ)
“Feels hot” because it releases energy to you
Endothermic
Reactions that absorb energy
Energy is a reactant in the reaction
N2 + O kJ → 2NO
 H is positive ( H= kJ)
“Feels cold” because it absorbs energy from you

5. D. Potential Energy Diagrams
Exothermic Reaction
Ea forward
Ea reverse
 H
Hreactant
HActivated Complex
Hproduct

6. D. Potential Energy Diagrams
Endothermic Reaction
1 Hreactant
2 Hactivated complex
3 Hproduct
4 Ea forward
5  H
6 Ea reverse

7. E. Collision Theory
A reaction occurs when particles collide with sufficient energy and proper alignment
More collisions, faster rate

8. F. Factors Affecting the Rate of Reactions Concentration
Increasing concentration increases number of molecules
Increases collisions
Increases rate
Temperature
Increasing temperature increases speed of molecules
Surface Area
Increasing surface area increases contact with reactants

9. Pressure Nature of Reactant Catalyst
Increase pressure decreases volume of a gas; less space
Increases collisions
Increases rate
Nature of Reactant
Ionic bonds break into pieces which increase concentration
Covalent bonds do not breakdown ; react slower
Catalyst
Changes the mechanism of the reaction
Lowers activation energy making reaction occur easier
More collisions can occur
Rate increases

10. II Equilibrium
Equilibrium is the balance between the rates of two opposing reactions
Example of two opposing reactions:
H2O(s) → H2O(l) melting
H2O(s) ← H2O(l) freezing
written in equilibrium notation
H2O(s) ↔ H2O(l) phase change
forward direction—melt
reverse direction—freeze
At equilibrium,
Rates are equal
Concentrations are constant

11. A. Le Chatelier’s Principle
systems that are at equilibrium are stable and want to remain at equilibrium
when equilibrium reactions are “stressed”, they will “shift” to establish equilibrium
stresses include:
concentration changes
temperature changes
pressure changes or volume changes (gases only)
add a catalyst

12. 1. Concentration Changes
When changing concentration, use the teeter-tauter technique
Given the reaction:
2 NO2 ↔ N2O4
What will happen to [N2O4 ] if [NO2 ] is increased?
([ ] represents concentration)
Tilt left
Shift right
Makes more N2O4

13. 2. Temperature Changes When changing temperature,
Use the teeter-tauter technique OR
If temp increases, shift away from heat
If temp decreases, shift toward heat
Given the reaction:
2 NO kJ ↔ N2O4
What will happen to [N2O4 ] if the temperature is increased?
Tilt left
Shift right
Makes more N2O4
Heat (300 kJ) is on the left
Equilibrium shifts away from the left side
Shifts right
N2O4 increases

14. 3. Pressure Changes Volume Changes
When pressure changes occur, only gases will be effected (g)
Count gases in the system
An increase in pressure (decrease in volume) causes a shift toward the smallest side of the reaction
Given the reaction:
2 NO2 (g) ↔ N2O4 (g)
What will happen to [N2O4 ] if the pressure is increased?
Count gases
2 gases left side……….1 gas right side
Shifts right (fewer gases)
Makes more N2O4

15. 4. Addition of a Catalyst
When a catalyst is added to a system at equilibrium, both forward and reverse reactions increase rate
There is no effect on the equilibrium
No shift will occur

16. Test yourself Given the reaction:
N2(g) + 3 H2 (g) ↔ 2 NH3 (g) + 80 kJ
What will happen to the concentration of NH3 if
H2 decreases?
Temperature increases?
Pressure decreases?
A catalyst is added?

17. III Enthalpy and Entropy
There are TWO factors which determine if a reaction will occur spontaneously or not
Enthalpy (ΔH)
The natural tendency is to change to a lower energy state
Exothermic direction is preferred
Entropy (ΔS)
Entropy measures randomness or disorder
Greater disorder (messy), higher entropy
(solid) lowest entropy
(liquid)
(aqueous)
(gas) highest entropy
High entropy is preferred

18. Spontaneous or Nonspontaneous?
C(s) + O2(g) ↔CO2(g) + 120kJ
∆H = -120kJ
Exothermic
∆ S solid/gas to gas only
Increase in entropy
Spontaneous

19. Spontaneous or Nonspontaneous?
N2 (g) + 4H2(g) + Cl2(g) + 93kJ ↔ 2NH4Cl(s)
∆H = +93kJ
Endothermic
∆ S gas to solid
Decrease in entropy
Nonspontaneous