1. CHEMICAL EQUILIBRIUM
Cato Maximilian Guldberg and his brother-in-law Peter Waage developed the Law of Mass Action

2. Chemical Equilibrium Reversible Reactions:
A chemical reaction in which the products
can react to re-form the reactants
Chemical Equilibrium:
When the rate of the forward reaction
equals the rate of the reverse reaction
and the concentration of products and
reactants remains unchanged
2HgO(s)  2Hg(l) + O2(g)
Arrows going both directions (  ) indicates equilibrium in a chemical equation

3. 2NO2(g)  2NO(g) + O2(g) Remember this from Chapter 12?
Why was it so important to measure reaction rate at the start of the reaction
(method of initial rates?)

4. 2NO2(g)  2NO(g) + O2(g)

5. jA + kB  lC + mD Law of Mass Action For the reaction:
Where K is the equilibrium constant, and is unitless

6. Product Favored Equilibrium
Large values for K signify the reaction is “product favored”
When equilibrium is achieved, most reactant has been converted to product

7. Reactant Favored Equilibrium
Small values for K signify the reaction is “reactant favored”
When equilibrium is achieved, very little reactant has been converted to product

8. Writing an Equilibrium Expression
Write the equilibrium expression for the reaction:
2NO2(g)  2NO(g) + O2(g)
K = ???

9. Example: Writing an Equilibrium Expression
Write the equilibrium expression for the reaction:

10. Conclusions about Equilibrium Expressions
The equilibrium expression for a reaction is the reciprocal for a reaction written in reverse
2NO2(g)  2NO(g) + O2(g)
2NO(g) + O2(g)  2NO2(g)

11. Conclusions about Equilibrium Expressions
When the balanced equation for a reaction is multiplied by a factor n, the equilibrium expression for the new reaction is the original expression, raised to the nth power.
2NO2(g)  2NO(g) + O2(g)
NO2(g)  NO(g) + ½O2(g)

12. Example: Calculating the Value of K
The following equilibrium concentrations were observed for the Haber process at 1270C.
a. Calculate the value of K
at 1270C for this reaction
b. Calculate the value of the equilibrium constant at 1270C for reaction:

13. Example: Calculating the Value of K
The following equilibrium concentrations were observed for the Haber process at 1270C.
c. Calculate the value of the equilibrium constant at 1270C for reaction given by the equation:

14. Equilibrium Expressions Involving Pressure
For the gas phase reaction:
3H2(g) + N2(g)  2NH3(g)

15. Equilibrium Expressions Involving Pressure
For the gas phase reaction:
3H2(g) + N2(g)  2NH3(g)

16. Example: Equilibrium Expressions Involving Pressure
The reaction for the formation of nitrosyl chloride:
Was studied at 250C. The pressures at equilibrium were found to be:
Calculate the value of Kp for this reaction at 250C.
b. Calculate the value of K for this reaction at 250C.

17. Heterogeneous Equilibria
The position of a heterogeneous equilibrium does not depend on the amounts of pure solids or liquids present
Write the equilibrium expression for the reaction:
PCl5(s)  PCl3(l) + Cl2(g)
Pure
solid
Pure
liquid

18. Equilibrium Expressions for Heterogeneous Equilibria
Write the expressions for K and Kp for the following process:
a. Solid phosphorus pentachloride decomposes to liquid phosphorus trichloride & chlorine gas.
b. Deep blue solid copper (II) sulfate pentahydrate is heated to drive off water vapor to from solid white copper (II) sulfate.

19. jA + kB  lC + mD The Reaction Quotient
For some time, t, when the system is not at equilibrium, the reaction quotient, Q takes the place of K, the equilibrium constant, in the law of mass action.
jA + kB  lC + mD

20. Significance of the Reaction Quotient
If Q = K, the system is at equilibrium
If Q > K, the system shifts to the left, consuming products and forming reactants until equilibrium is achieved
If Q < K, the system shifts to the right, consuming reactants and forming products until equilibrium is achieved

21. Using the Reaction Quotient
For the synthesis of ammonia at 5000C, the equilibrium constant is 6.0x10-2 Predict the direction in which the system will shift to reach equilibrium in each of the following cases:
a. [NH3]0 = 1.0x10-3 M; [N2]0 = 1.0x10-5 M; [H2]0 = 2.0x10-3 M
b. [NH3]0 = 2.0x10-4 M; [N2]0 = 1.5x10-5 M; [H2]0 = 3.54x10-1 M
c. [NH3]0 = 1.0x10-4 M; [N2]0 = 5.0 M; [H2]0 = 1.0x10-2 M

22. Solving for Equilibrium Concentration
Consider this reaction at some temperature:
H2O(g) + CO(g)  H2(g) + CO2(g) K = 2.0
Assume you start with 8 molecules of H2O and 6 molecules of CO. How many molecules of H2O, CO, H2, and CO2 are present at equilibrium?
Here, we learn about “ICE” – the most important problem solving technique in the second semester. You will use it for the next 4 chapters!

23. Solving for Equilibrium Concentration
H2O(g) + CO(g)  H2(g) + CO2(g) K = 2.0
Step #1: We write the law of mass action for the reaction:

24. Solving for Equilibrium Concentration
Step #2: We “ICE” the problem, beginning with the Initial concentrations
H2O(g) + CO(g)  H2(g) + CO2(g)
Initial:
Change:
Equilibrium: 8 6 -x -x +x +x
8-x
6-x x x

25. Solving for Equilibrium Concentration
Step #3: We plug equilibrium concentrations into our equilibrium expression, and solve for x
H2O(g) + CO(g)  H2(g) + CO2(g)
Equilibrium:
8-x
6-x x
x = 4

26. Solving for Equilibrium Concentration
Step #4: Substitute x into our equilibrium concentrations to find the actual concentrations
H2O(g) + CO(g)  H2(g) + CO2(g)
Equilibrium:
8-x
6-x x
x = 4
Equilibrium:
8-4=4
6-4=2 4

27. Calculating Equilibrium Pressure
Dinitrogen tetroxide in its liquid state was used as one of the fuels on the lunar landing for the NASA Apollo missions. In the gas phase it decomposes to gaseous nitrogen dioxide:
Consider an experiment in which gaseous was placed in a flask & allowed to come to equilibrium at a temperature where Kp= At equilibrium, the pressure of was found to be at 2.71 atm. Calculate the equilibrium pressure of

28. Calculating Equilibrium Concentrations
Assume that the reaction for the formation of gaseous hydrogen fluoride from hydrogen & fluorine has an equilibrium constant 0f 1.15x102 at a certain temperature. In an experiment, mol of each component was added to a L flask. Calculate the equilibrium concentrations of all species.

29. Solving Equilibrium Problems
1. Write a balanced equation.
2. Write the equilibrium expression using the law of mass action.
3. If given initial concentrations, use Q (reaction quotient) to predict the direction the reaction will proceed.
4. ICE the problem.
I – initial concentration
C – Change in concentration of reactants & products (remember stoichiometry here!)
E – Equilibrium concentrations
5. Substitute the equilibrium concentrations into the equlibrium expression and solve for the unknown (x).
6. Check your calculated equilibrium [ ] by making sure they give the correct value of K.

30. Calculating Equilibrium Pressures
Assume that the reaction for the formation of gaseous hydrogen iodide from hydrogen gas & iodine vapor has an equilibrium constant of 1.00 x 102 at a certain temperature. Suppose HI at x atm, H2 at x atm, & I2 at x atm are mixed in a L flask. Calculate the equilibrium concentrations of all species.

31. Calculating Equilibrium Pressures
K=1.00 x 102
HI=5.000 x atm, H2=1.000 x atm, I2=5.000 x atm

32. Treating systems that have small equilibrium constants
We can make some approximations when the equilibrium constant is much smaller than the concentrations of the reactants & products.
You must check to see if your approximation is valid using the “5% rule”
You must note your approximation in the problem
As a general rule of thumb – if the equilibrium constant is smaller than the concentrations by a magnitude of 103 or more – an approximation can be made.

33. Calculating Equilibrium Concentrations with small equilibrium constants
Gaseous NOCl decomposes to form the gases NO & Cl2. At 350C K = 1.6x In an experiment in which 1.0 mol of NOCl is placed in an empty 2.0 L flask, what the equilibrium concentrations of all species?

34. Le Chatelier’s Principle

35. LeChatelier’s Principle
When a system at
equilibrium is placed under
stress, the system will
undergo a change in such
a way as to relieve that
stress and restore a
state of equilibrium.
Henry Le Chatelier

36. Le Chatelier Translated:
When you take something away from a system at equilibrium, the system shifts in such a way as to replace some of what you’ve taken away.
When you add something to a system at equilibrium, the system shifts in such a way as to use up some of what you’ve added.

37. LeChatelier Example #1
A closed container of ice and water is at equilibrium. Then, the temperature is raised.
Ice + Energy  Water
The system temporarily shifts to the _______ to restore equilibrium.
right

38. LeChatelier Example #2
A closed container of N2O4 and NO2 is at equilibrium. NO2 is added to the container.
N2O4 (g) + Energy  2 NO2 (g)
The system temporarily shifts to the _______ to restore equilibrium.
left

39. LeChatelier Example #3
A closed container of water and its vapor is at equilibrium. Vapor is removed from the system.
water + Energy  vapor
The system temporarily shifts to the _______ to restore equilibrium.
right

40. LeChatelier Example #4
A closed container of N2O4 and NO2 is at equilibrium. The pressure is increased.
N2O4 (g) + Energy  2 NO2 (g)
The system temporarily shifts to the _______ to restore equilibrium, because there are fewer moles of gas on that side of the equation.
left