1. Reaction Rates & Equilibrium
Unit 12 - Chapter 18 1

2. Reaction Rate
Reaction rate – how fast reactants disappear and how fast product appears

3. A B
time
rate = -
D[A] Dt
rate =
D[B] Dt
13.1

4. Reaction Rate Reaction Rate = ∆ [A] ∆ t Example:
CO(g) + NO2(g)  CO2(g) + NO(g)
- at t = 4.0 min, [CO2] = .12 mol/L
- at t = 8.0 min, [CO2] = .24 mol/L
- reaction rate = .24 mol/L mol/L
8.0 min – 4.0 min
= mol/L . min
Unit for reaction rate = conc. with some time unit
Products have a (+) rate
Reactants have a (-) rate

5. Collision Theory of Kinetics
Kinetics is the study of the factors that affect the speed of a reaction and the mechanism by which a reaction proceeds.
In order for a reaction to take place, the reacting molecules must collide into each other.
Once molecules collide they may react together or they may not, depending on two factors -
Whether the collision has enough energy to "break the bonds holding reactant molecules together";
Whether the reacting molecules collide in the proper orientation for new bonds to form. 2

7. Effective Collisions
Collisions in which these two conditions are met (and therefore the reaction occurs) are called effective collisions.
The higher the frequency of effective collisions the faster the reaction rate.
When two molecules have an effective collision, a temporary, high energy (unstable) chemical species is formed - called an activated complex
It is a transition state between reactant and product
It has a very short lifetime (10-13 s)
Has to form for product to be formed 3

8. Activated Complex
The difference in potential energy between the reactant molecules and the activated complex is called the activation energy, Ea
This is the minimum amount of energy that particles must have in order to react.
The larger the activation energy, the slower the reaction
The energy to overcome the activation energy comes from the kinetic energy of the collision being converted into potential energy, or from energy available in the environment, i.e. heat.
Different reactions have different activated complexes and therefore different activation energies 4

9. Energy Diagram Energy of products is lower than energy of reactants
What is this called?
Energy of products is lower than energy of reactants
energy lost, exothermic, -∆H
Energy of products is higher than energy of reactants
energy gained, endothermic, +∆H

10. Factors Affecting Reaction Rate
Nature of the Reactants
Cl2(g) + CH4(g)  CH3Cl(g) + HCl(g)
Cl2  Cl + Cl (fast)
Cl + CH4  CH3Cl + H (slow)
H + Cl  HCl (very fast)
individual steps = elementary steps
all steps together = reaction mechanism
the slowest step determines the rate of the reaction
called the rate determining step
Intermediates – product in one step, reactant in another

11. Factors Affecting Reaction Rate
Concentration
The larger the concentration of reactant molecules, the faster the reaction will go.
Increases the frequency of reactant molecule collisions
3. Particle Size (Surface Area)
more particles on the surface = more particles available for collisions
more collisions = more act. complex = more product
smaller particles give you more surface area

12. Factors Affecting Reaction Rate
4. Agitation
this puts more liquid/gas particles in contact with the solid = ↑ collisions = ↑ act. complex = ↑ product
5. Pressure
↑ pressure by ↓ volume – puts particles closer together = ↑ collisions = ↑ act. complex = ↑ product
All of these factors are similar, in terms of explanation, to concentration!!!

13. Factors Affecting Reaction Rate
6. Temperature
most effective at speeding up a reaction
↑ temp. = ↑ KE (particles moving faster)
particles move faster leading to more collisions
the collisions are also harder
these harder collisions contain the needed energy to overcome the Ea
therefore the reaction rate will increase

14. Factors Affecting Reaction Rate
7. Catalyst
substance that speeds up a reaction, but isn’t used up in the reaction
provides a “different pathway” that requires lower Ea
lower Ea = more collisions having the proper amount of energy = ↑ act. complex = ↑ product

15. Reaction Dynamics
If the products of a reaction are removed from the system as they are made, then a chemical reaction will proceed until the limiting reactants are used up.
However, if the products are allowed to accumulate; they will start reacting together to form the original reactants - called the reverse reaction.
We show this reverse reaction by using a double arrow (H2(g) + I2(g)  2HI(g)) 10

16. Reaction Dynamics
The forward reaction slows down as the amounts of reactants decreases because the reactant concentrations are decreasing
At the same time the reverse reaction speeds up as the concentration of the products increases.
Eventually the forward reaction is using reactants and making products as fast as the reverse reaction is using products and making reactants. This is called chemical equilibrium.
rateforward = ratereverse
Note: This equilibrium is dynamic 11

17. Chemical Equilibrium
Dynamic Equilibrium can only be reached in a closed system!!
When a system reaches equilibrium, the amounts of reactants and products in the system stays constant
the forward and reverse reactions still continue, but because they go at the same rate the amounts of materials don't change.
There is a mathematical relationship between the amounts of reactants and products at equilibrium 12

18. Equilibrium Expression=
[C]c[D]d
[A]a[B]b
Keq
aA + bB Û cC + dD
Capital letters (A,B,C,D) – reactants or products
Lowercase letter (a,b,c,d) – coefficients from the equation
NOTE – products on top, reactants on bottom
In this expression, Keq is a number called the equilibrium constant.
ratio of product concentration to reactant concentration at equilibrium
Do not include solids or liquids, only solutions and gases
The value of Keq depends on temp. of the reaction – if temp. changes then the value of Keq changes. 13

19. So what does this Keq value tell us???
Example – Determine the value of the Equilibrium Constant for the Reaction 2 SO2(g) + O2(g) Û 2 SO3(g)
Determine the Equilibrium Expression
Plug the equilibrium concentrations into to Equilibrium Expression
Solve the Equation
3.50
3.00
SO3
1.25
1.50 O2
2.00
SO2
[Equilibrium]
[Initial]
Chemical
So what does this Keq value tell us??? 16

20. Position of Equilibrium
The size of the equilibrium constant shows whether products or reactants are favored at equilibrium.
Keq > 1, products are favored at equilibrium
Keq < 1, reactants are favored at equilibrium 14

21. Example – If the value of the Equilibrium Constant for the Reaction 2 SO2 + O2 Û 2 SO3 is 4.36, Determine the Equilibrium Concentration of SO3
Determine the Equilibrium Expression
Plug the equilibrium concentrations and Equilibrium Constant into the Equilibrium Expression
Solve the Equation ?
3.00
SO3
1.25
1.50 O2
2.00
SO2
[Equilibrium]
[Initial]
Chemical 21

22. More Equilibrium Practice
1. Write the equilibrium constant expression for the following reaction.
3H2(g) + N2(g) ↔ 2NH3(g)
2. An analysis of an equilibrium mixture for this reaction in a 1.0 L flask at 300oC gave the following results: 0.15 mol H2, 0.25 mol N2 and 0.10 mol NH3. Calculate the Keq for this reaction.
3. 2BrCl(g) ↔ Cl2(g) + Br2(g)
The equilibrium constant for this reaction is The equilibrium mixture contains 4.00 mol Cl2 and 4.00 moles of Br2. How many moles of BrCl are present?

23. Le Châtelier’s Principle
Le Châtelier's Principle guides us in predicting the effect various changes have on the position of equilibrium
it says that if stress is applied to a system in dynamic equilibrium, the system will change to relieve the stress.
The position of equilibrium moves to counteract the change.
Three common stressors:
Concentration
Temperature
Pressure 17

24. Concentration Changes and Le Châtelier’s Principle
A + B ↔ C D
Adding a reactant – equilibrium shifts right
Removing a reactant – equilibrium shifts left
Adding a product – equilibrium shifts left
Removing a product – equilibrium shifts right 18

25. Changing Pressure and Le Châtelier’s Principle
Only affects a reaction involving gases with an unequal number of mole of reactants & products.
Increasing the pressure on the system causes the position of equilibrium to shift toward the side of the reaction with the fewer gas molecules
Decreasing pressure causes a shift toward the side with more gas molecules
Example
3H2(g) + N2(g)  2NH3(g) kJ
↑ Pressure – shifts to the right
↓ Pressure – shift to the left 19

26. Changing Temperature and Le Châtelier’s Principle
Increasing the temperature causes the reaction to shift away from the heat.
For exothermic reactions -
Think of heat as a product of the reaction
Therefore shift the position of equilibrium toward the reactant side
For endothermic reactions -
Think of heat as a reactant
The position of equilibrium will shift toward the products
Cooling an exothermic or endothermic reaction will have the opposite effects. 20

27. Examples – Le Chatelier’s Principle
What effect do the following changes have on the equilibrium position for the following reaction?
1. PCl5(g) + heat ↔ PCl3(g) + Cl2(g)
a. addition of Cl2
b. increase in pressure
c. removal of heat
d. removal of PCl3 as formed
2. C(s) + H2O(g) + heat ↔ CO(g) + H2(g)
a. Lowering the temperature
b. Increasing the pressure
c. Removal of H2 as formed

28. Examples – Le Chatelier’s Principle
3. At 425 K –
Fe3O4(s) + 4H2(g) ↔ 3Fe(s) + 4H2O(g)
How would the equilibrium concentration of H2O be affected by the following:
a. Adding more H2
b. Adding more Fe(s)
c. Decreasing the pressure
d. Adding a catalyst