1. Equilibrium -OR- The Heart and Soul of AP Chemistry
The Bane Of My Existence
(if I’m an AP Chem Student)

2. Reactions are reversible
A + B C + D ( forward)
C + D A + B (reverse)
Initially there is only A and B so only the forward reaction is possible
As C and D build up, the reverse reaction speeds up while the forward reaction slows down.
Eventually the rates are equal (not necessarily the concentrations)

3. Forward Reaction
Reaction Rate
Equilibrium
Reverse reaction
Time

4. What is equal at Equilibrium?
Rates are equal
Concentrations are not.
Rates are determined by concentrations and activation energy.
The concentrations do not change at equilibrium.
or if the reaction is verrrry slooooow.

5. Law of Mass Action For any reaction jA + kB lC + mD
K = [C]l[D]m PRODUCTSpower [A]j[B]k REACTANTSpower
K is called the equilibrium constant.
is how we indicate a reversible reaction

6. Playing with K If we write the reaction in reverse. lC + mD jA + kB
Then the new equilibrium constant is
K’ = [A]j[B]k = 1/K [C]l[D]m

7. K’ = [A]nj[B]nk = ([A] j[B]k)n = Kn [C]nl[D]nm ([C]l[D]m)n
Playing with K
If we multiply the equation by a constant
njA + nkB nlC + nmD
Then the equilibrium constant is
K’ = [A]nj[B]nk = ([A] j[B]k)n = Kn [C]nl[D]nm ([C]l[D]m)n

8. The units for K
Are determined by the various powers and units of concentrations.
They depend on the reaction.
Rarely used in problems – sometimes show up in MC section

9. K is CONSTANT Temperature affects rate.
for any given temperature.
Temperature affects rate.
The equilibrium concentrations don’t have to be the same, only K.
Equilibrium position is a set of concentrations at equilibrium.
There are an unlimited number.

10. There can be only one. . . (for each Temperature)
Equilibrium Constant
There can be only one. . .
(for each Temperature)

11. Equilibrium and Pressure
Some reactions are gaseous
PV = nRT
P = (n/V)RT
Let C beconcentration in moles/Liter
P = CRT

12. Equilibrium and Pressure
2SO2(g) + O2(g) SO3(g)
Kp = (PSO3) (PSO2)2 (PO2)
Kc = [SO3] [SO2]2 [O2]

13. General Equation jA + kB lC + mD Kp = Kc (RT)Dn
Dn - Change in moles of gas (products-reactants)

14. Solving Equilibrium Problems
Type 1

15. The Reaction Quotient
Tells you the directing the reaction will go to reach equilibrium
Calculated the same as the equilibrium constant, but for a system not at equilibrium
Q = [Products]coefficient [Reactants] coefficient
Compare value to equilibrium constant

16. What Q tells us If Q<K If Q>K If Q=K system is at equilibrium
Not enough products
Shift to right
If Q>K
Too many products
Shift to left
If Q=K system is at equilibrium

17. Example for the reaction 2NOCl(g) 2NO(g) + Cl2(g)
K = 1.55 x 10-5 M at 35ºC
In an experiment mol NOCl, mol NO(g) and mol Cl2 are mixed in 2.0 L flask.
Which direction will the reaction proceed to reach equilibrium?

18. Solving Equilibrium Problems
Given the starting concentrations and one equilibrium concentration.
Use stoichiometry to figure out other concentrations and K.
Learn to create a table of initial and final conditions.

19. Consider the following reaction at 600oC
2SO2(g) + O2(g) SO3(g)
In a certain experiment 2.00 mol of SO2, 1.50 mol of O2 and 3.00 mol of SO3 were placed in a 1.00 L flask. At equilibrium 3.50 mol of SO3 were found to be present. Calculate the equilibrium concentrations of O2 and SO2, K and KP

20. I C E 2SO2(g) O2(g) 2SO3 2.00mol/1.0L 1.50mol/1.00L 3.00mol/1.00L -2x
0.50mo/L = 2x
0.25mol/L = x
[SO2]eq = 2.00mol/L - 2(0.25mol/L) = 1.50mol/L
[O2]eq = 1.50mol/L – 0.25mol/L = 1.25mol/L

21. What if you’re not given equilibrium concentration?
The size of K will determine what approach to take.
Very large (K>100) or very small (K<0.01)
Allows us to make simplifying assumptions.

22. For LARGE K values:
IF K>>1, more product will be formed.
Assumption: since K is large reaction will go to completion.
Set up ICE Chart
“-x” on reactant side, “+x” on product side
Use 5% rule to drop X from denominator

23. For SMALL K values:
Reaction will be product favored
Set up ICE Chart
“+x” on reactant side, “-x” on product side
Use 5% rule to drop X from denominator

24. Checking the assumption THE 5% RULE
The rule of thumb is that if the value of X is less than 5% of all the other concentrations, our assumption was valid.
If not we would have had to use the quadratic equation
More on this later.
Our assumption was valid.

25. MID-RANGE K 0.01<K<10 Determine direction of shift by value of K
Can’t use 5% rule, so we will have to solve the quadratic (Yeah, it sucks)
Good News!
A) you can get a quadratic solver for your TI calculator to use on first FR section
B) quadratics RARELY appear on AP test

26. Practice Problem What are the equilibrium concentrations.
For the reaction 2ClO(g) Cl2 (g) + O2 (g)
K = 6.4 x 10-2
In an experiment mol ClO(g), 1.00 mol O2 and 1.00 x 10-2 mol Cl2 are mixed in a 4.00 L container.
What are the equilibrium concentrations.

27. Problems Involving Pressure
Solved exactly the same, with same rules for choosing X depending on KP

28. Le Chatelier’s Principle
If a stress is applied to a system at equilibrium, the position of the equilibrium will shift to reduce the stress.
According to LeChatelier’s Principle. . .
This is a label applied to describe a phenomena, NOT an explanation

29. Change amounts of reactants and/or products
Adding product makes Q>K
Removing reactant makes Q>K
Adding reactant makes Q<K
Removing product makes Q<K
Determine the effect on Q, will tell you the direction of shift

30. Change Pressure By changing volume
System will move in the direction that has the least moles of gas.
Because partial pressures (and concentrations) change a new equilibrium must be reached.
System tries to minimize the moles of gas.

31. Change in Pressure By adding an inert gas
Partial pressures of reactants and product are not changed
No effect on equilibrium position

32. Change in Temperature
Affects the rates of both the forward and reverse reactions.
Doesn’t just change the equilibrium position, changes the equilibrium constant.
The direction of the shift depends on whether it is exo- or endothermic

33. Exothermic DH<0 Releases heat Think of heat as a product
Raising temperature push toward reactants.
Shifts to left.

34. Endothermic DH>0 Produces heat Think of heat as a reactant
Raising temperature push toward products.
Shifts to right.