1. Rates of Reaction and Chemical Equilibrium
Chemistry
Ms. Piela

2. Rates of Reaction
A rate is any change that occurs within an interval of time
Collision theory states that particles will react to form products when they collide
This requires a certain amount of kinetic energy
As an example, imagine throwing clay balls at one another
Without enough force, they will ricochet off from one another. When they have enough kinetic energy, they will smash and fuse into something entirely new.
This is the same in chemical reactions

3. Activation Energy
Activation energy (Ea) is the minimum amount of energy required for a chemical reaction to occur
Think of activation energy like a barrier

4. Activation Energy Diagram
Diagram can be used to determine in reaction is exothermic (- ∆H) or endothermic (+∆H)
Activated complex (or transition state) is the arrangement of atoms at the peak of the Ea barrier
Unstable and only appears briefly

5. Activation Energy Diagram

6. Factors Affecting Rates of Reactions
Four Major Factors
Temperature
Concentration
Particle Size
Catalysts

7. Factors Affecting Rates of Reaction
Temperature
Generally, an increase in temperature will increase the rate of a reaction, and vice versa
Explanation:
Increasing temperature speeds up particles, increasing kinetic energy
This allows a reaction to overcome its activation energy
Example
Popcorn

8. Factors Affecting Rates of Reaction
Concentration
Increasing the number of particles (thereby increasing concentration) will increase the reaction rate
Explanation:
More particles will increase the collision frequency
More collisions will mean more chances for a reaction to occur
Example
Wooden splint in pure oxygen

9. Factors Affecting Rates of Reaction
Particle Size (aka Surface Area)
The smaller the particle, the larger the surface area for a given mass of particles
An increase in surface area will increase the rate of reaction
Explanation:
With more exposed area to react, this increases the rate of reaction
Example
Fine Dust Explosions (1902 Chicago Fire)
Lycopodium dust explosion

10. Factors Affecting Rates of Reaction
Catalysts
Any substance that increases the rate of a reaction without being consumed
Catalysts allow a reaction to occur at lower activation energy
Enzymes are biological catalysts

11. Activation Energy Diagram

12. Factors Affecting Rates of Reaction

13. Review of Factors Affecting Rates
Temperature
Increase causes higher kinetic energy
Concentration
Higher concentration causes more collisions
Particle Size
More surface area leads to more exposed spaces to react
Catalysts
Substances that increase the rates of reaction without being spent

14. Chemical Equilibrium
Reversible reactions are reactions that simultaneously occur in both directions
Chemical equilibrium is the state in which the forward and reverse reactions take place at the same rate
Sample Reaction:
2 SO2 (g) + O2 (g)  2 SO3 (g)

15. Chemical Equilibrium Differences between rates and concentration
One doesn’t imply the other!
Although rates of forward and reverse reactions are equal at chemical equilibrium, the concentrations are not necessarily the same
Example
A  B
1% %
If A reacts to give B and the equilibrium mixture contains 99% B and only 1% A, the formation of B is said to be favored

16. Effect of Catalyst on Equilibrium
Although catalysts speed up a reaction rate, they will not affect the concentration of reactants and products at equilibrium
Catalysts can only decrease the time required to reach equilibrium
Affects speed, not concentration! Meep meep!

17. Le Châtelier's Principle
When a stress is applied to a system at equilibrium, the equilibrium will shift in order to relieve the stress

18. Factors Affecting Equilibrium
Three Major Factors
Concentration
Temperature
Pressure

19. Le Châtelier's Principle
Concentration
Changing the amount, or concentration, of any reactant or product in a system at equilibrium disturbs that equilibrium
Example equilibrium:
2 H2CO3 (g)  2 H2O (l) + 2 CO2 (g)
Adding CO2
Shifts equilibrium to the left
Concentration of H2O will decrease
Concentration of H2CO3 will increase

20. Le Châtelier's Principle
Example equilibrium:
2 H2CO3 (g)  2 H2O (l) + 2 CO2 (g)
Removing CO2
Shifts equilibrium to the right
Concentration of H2O will increase
Concentration of H2CO3 will decrease

21. Le Châtelier's Principle
Temperature
Increasing the temperature causes an equilibrium position of a reaction to shift in the direction that will absorb the heat.
Consider heat as if it were a part of the reaction
2 SO2 (g) + O2 (g)  2 SO3 (g) + heat
Adding heat
Shifts equilibrium to the left.
Concentration of SO2 and O2 will increase.
Concentration of SO3 will decrease.

22. Le Châtelier’s Principle
2 SO2 (g) + O2 (g)  2 SO3 (g) + heat
Removing heat
Shifts equilibrium to the right.
Concentration of SO2 and O2 will decrease.
Concentration of SO3 will increase.

23. Le Châtelier's Principle
Pressure
Changing the pressure only affects an equilibrium with an uneven number of reactants and products
Increasing pressure will shift an equilibrium in the direction of the least number of moles
Be aware that a change in volume also affects pressure
Reducing volume will increase the pressure, and vice versa

24. Le Châtelier's Principle
Example Equilibrium:
CO + 3 H2 (g)  CH4 (g) + H2O (g)
Increasing pressure (reducing volume) will shift the equilibrium to the…
right
Decreasing pressure (increasing volume) will shift the equilibrium to the…
left

25. Haber Bosch Process & Le Châtelier’s Principle

26. Equilibrium Constant Expressions
A mathematical relationship exists between concentration of the reactants and products once equilibrium has been reached that is independent of the initial concentration of the participants
Given the symbol Keq
Can be used to determine whether products or reactants are favored in an equilibrium

27. Equilibrium Constant Expressions

28. Steps for Determining Keq
Make sure the chemical reaction is balanced and at equilibrium
Place the product concentrations in the numerator, and the reactant concentrations in the denominator
[products]
[reactants]
Concentrations of any solid or liquid is left out because the concentrations never change
Water is always omitted in aqueous solution. Water is almost constant during the reaction
Keq =

29. Steps for Determining Keq
To complete the expression, raise each substance’s concentration to the power equal to the substance’s coefficient in the balanced chemical equation
For a general reaction: aA+ bB  cC + dD
The equilibrium constant expression can be written as
The equilibrium constant is unitless
[C]c[D]d [A]a[B]b
Keq =

30. H2CO3 (aq) + H2O (l)  HCO3- (aq) + H3O+ (aq)
Example #1
An aqueous solution of carbonic acid reacts to reach equilibrium described below:
H2CO3 (aq) + H2O (l)  HCO3- (aq) + H3O+ (aq)
Write the equilibrium expression, Keq, for this reaction

31. N2O4 (colorless)  2 NO2 (brown)
Example #2
The brown gas, NO2, is in equilibrium with the colorless gas N2O5. Write the equilibrium expression for the following reaction:
N2O4 (colorless)  2 NO2 (brown)

32. Practice Problems
Practice #1
Practice #2

33. H2CO3 (aq) + H2O (l)  HCO3- (aq) + H3O+ (aq)
Calculating using Keq
Example #1
H2CO3 (aq) + H2O (l)  HCO3- (aq) + H3O+ (aq)
The solution contains the following solute concentrations: H2CO3; 3.3 x 10-2 M; HCO3- ion, 1.19 x 10-4 M; and H3O+, 1.19 x 10-4 M. Determine the Keq expression, and the value of Keq

34. Example #1

35. NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
Example #2
Keq for the equilibrium below is 1.8 x 10-5 at a temperature of 25 °C. Calculate [NH4+] when [NH3] = 6.82 x 10-3 M and [OH-] = 3.45 x 10-3 M.
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)

36. Example #2

37. Practice Problems
Problem #1
Problem #2

38. Magnitude of Keq The meaning of Keq
If Keq is very large, products will be favored
This implies the numerator is large in the Keq expression compared to denominator
If Keq is very small, reactants will be favored
This implies the denominator is large in the Keq expression compared to numerator
If Keq is close to 1, then roughly equal amounts of reactants and products are present at equilibrium