1. Section 8.1—Equilibrium
What is equilibrium?

2. Reversible Reactions
Reversible Reactions – A chemical reaction that can proceed in both directions (represented by a “”)
Molecules A & B collide to produce molecules C & D
A + B C + D
But in some reactions, molecules C & D can collide and produce A & B again
A + B C + D
Reversible reactions are shown with a double arrow
A + B C + D

3. Collision Theory & Reaction Rate
Recall that molecules MUST collide with proper energy and orientation before they can react!!!
The more molecules there are, the more often these successful collisions will occur—Reaction rate will be faster.
At first the concentration of reactants is high—the forward reaction has a high reaction rate
As reactants start to turn into products, the reaction rate lowers (fewer molecules)
Reaction rate
Time
Forward reaction

4. Collision Theory & Reaction Rate
But in reversible reactions, once there are products the reverse reaction can occur as well.
At first the concentration of products is very low, the reaction rate of the reverse reaction is low
As products begin to form, the reverse reaction rate increases
Reaction rate
Time
Reverse reaction

5. Putting Forward & Reverse Together
Reaction rate
Time
Forward reaction
Reaction rate
Time
Reverse reaction
Reaction rate
Time
Forward reaction
Reverse reaction
Equilibrium is reached

6. Equilibrium
Equilibrium – When the rate of the forward reaction equals the rate of the reverse reaction.

7. Establishing equilibrium
It takes time to establish equilibrium
Reactants  Products
At first, there are only reactants present.
Only the forward reaction is possible.

8. Establishing equilibrium
It takes time to establish equilibrium
When equilibrium is established, the number of products and reactants doesn’t change…but the reaction keeps going.
Reactants  Products
But once there are products as well, they can begin to reform reactants. The reverse reaction becomes possible.
Once the rate of the forward and reverse process are equal, it is at equilibrium.

9. Dynamic Equilibrium
Dynamic Equilibrium – The reaction continues to proceed in both directions, but at the same rate.
The number of products and reactants no longer change, it may look as thought the reaction has stopped…
But the reactions continues!

11. Section 8.2—Equilibrium Constant
How can we describe a reaction at equilibrium?

12. Equilibrium Constant Expression
Equilibrium Constant Expression – Equation showing the ratio of the concentrations of products to reactants at equilibrium
We use brackets, [ ], to symbolize concentration!!!

15. Writing Equilibrium Constant Expressions
Write the product of the product concentrations on the top—take each one to a power of the coefficient from the balanced equation. 1
Write the product of the reactant concentrations on the bottom—also take each to the power of the balanced equation coefficient. 2
Example:
Write the equilibrium constant expression for the following:
2 H2 (g) + O2 (g)  2 H2O (g)

17. Writing Equilibrium Constant Expressions
Write the product of the product concentrations on the top—take each one to a power of the coefficient from the balanced equation. 1
Write the product of the reactant concentrations on the bottom—also take each to the power of the balanced equation coefficient. 2
Example:
Write the equilibrium constant expression for the following:
2 H2 (g) + O2 (g)  2 H2O (g)
[H2O] 2
K = 2
[O2]
[H2]

18. Heterogeneous Equilibrium
Homogeneous Equilibrium – All of the species are the same state of matter
2 H2 (g) + O2 (g)  2 H2O (g)
Heterogeneous Equilibrium – There are at least 2 states of matter
2 H2 (g) + O2 (g)  2 H2O (l)

19. Concentrations of Solids and Liquids
Solids and Liquids are PURE – they are not described using concentration terms.
Concentration means a part (solute) in the whole (solvent) If a substance is pure it can’t be a part of a whole 
THEREFORE – WE DO NOT INCLUDE SOLIDS OR LIQUIDS IN THE EXPRESSION!!!!

20. “K” Expressions with Solids or Liquids
If the “concentration” of a pure solid or liquid is constant, then it will not change during equilibrium and it is not written in the “K” expression.
2 H2 (g) + O2 (g)  2 H2O (g)
2 H2 (g) + O2 (g)  2 H2O (l)
H2O is not included in this “K” expression because it’s a liquid.
Only gases and aqueous solutions are included in “K” expressions!

21. Example #1—Writing K expression
Write the equilibrium constant expression for
Fe2O3 (s) + 3 H2 (g)  2 Fe (s) + 3 H2O (g)

22. Example #1—Writing K expression
Write the equilibrium constant expression for
Fe2O3 (s) + 3 H2 (g)  2 Fe (s) + 3 H2O (g)
Fe2O3 and Fe were not included in the K expression as they are solids!

23. Equilibrium Constant
Equilibrium Constant (K)– The number calculated from the equilibrium constant expression
“K” is different for every reaction at every temperature!

24. Example #2—Calculating K
Solve for the equilibrium constant for
Fe2O3 (s) + 3 H2 (g)  2 Fe (s) + 3 H2O (g)
If at equilibrium [H2] = 0.45 M and [H2O] = 0.18 M

25. Example #2—Calculating K
Solve for the equilibrium constant for
Fe2O3 (s) + 3 H2 (g)  2 Fe (s) + 3 H2O (g)
If at equilibrium [H2] = 0.45 M and [H2O] = 0.18 M
K = 0.064
Most instructors and textbooks do not require units for “K” as each one would be different

26. Meaning of Equilibrium Constant
In general….
[Products]
If K >1 or large…
[Reactants]
There is a much larger ratio of products to reactants at equilibrium
The reaction is said to “lie to the right” (products are on the right)
If K < 1 or small….
[Products]
[Reactants]
There is a much smaller ratio of products to reactants at equilibrium
The reaction is said to “lie to the left”
If K = 1 it means we have similar amounts of reactants and products at equililbrium.

27. Write the equilibrium constant expression for
Let’s Practice #1
Write the equilibrium constant expression for
N2 (g) + O2 (g)  2 NO (g)

28. Write the equilibrium constant expression for
Let’s Practice #1
Write the equilibrium constant expression for
N2 (g) + O2 (g)  2 NO (g)

29. If the equilibrium constant for
Let’s Practice #2
If the equilibrium constant for
N2 (g) + O2 (g)  2 NO (g) is 1.24 × 10-4, what can be said in general about this reaction at equilibrium?

30. If the equilibrium constant for
Let’s Practice #2
If the equilibrium constant for
N2 (g) + O2 (g)  2 NO (g) is 1.24 × 10-4, what can be said in general about this reaction at equilibrium?
The equilibrium constant is very small, so at equilibrium the concentration of products is much lower than reactants. The reaction lies to the left.

31. If the equilibrium constant for
Let’s Practice #3
If the equilibrium constant for
N2 (g) + O2 (g)  2 NO (g) is 1.24 × 10-4 and the equilibrium concentration of [N2] = M and [O2] = M, what is the equilibrium concentration of NO?

32. If the equilibrium constant for
Let’s Practice #3
If the equilibrium constant for
N2 (g) + O2 (g)  2 NO (g) is 1.24 × 10-4 and the equilibrium concentration of [N2] = M and [O2] = M, what is the equilibrium concentration of NO?
[NO]eq = M

33. Section 8.4—Le Chatelier’s Principle
How can we push a reaction to make more products?

34. Le Chatelier’s Principle
Le Chatelier’s Principle – when a stress is applied to a system in equilibrium – it will shift to relieve the stress.

35. What kinds of stresses can be put on a system in equilibrium?
Adding or removing a reactant or product (i.e. changing the concentration)
Increasing or decreasing the pressure
Adding or removing heat

37. Changes in Pressure – This only matters if you have GASES in the Reaction!
Pressure increases – In the reaction the pressure comes from the gases!
Reaction shifts to the side with least moles of gas to decrease pressure
Decrease volume
What would happen if we decreased the volume for this system?
2 SO2 (g) + O2 (g)  2 SO3 (g)
It will shift to the side that decreases the pressure – The left has 3 moles of gas the right has 2 moles of gas so the shift will go Right!
Reactions shifts to the side with the most moles of gas to increase pressure
Pressure decreases
Increase volume
What would happen if we increased the volume for this system?
2 SO2 (g) + O2 (g)  2 SO3 (g)
It will shift to the side that increases the pressure – Left

39. Endo & Exothermic
Endothermic Reaction – The reaction takes in energy…it absorbs energy and the products now have more energy than the reactants
Energy is a reactant in the reaction
Exothermic Reaction – The reaction gives off energy…it releases energy and the products have less energy than the reactants
Energy is a product in the reaction

40. Changing temperature—Endothermic
Reaction shifts to right
(get rid of extra reactants and make more products)
Increase temperature of endothermic reaction
Increasing a reactant
181kJ + 2HgO ↔ 2 Hg + O2
If we increase the temperature on this system it is as if we are adding a reactant so it will shift away from the heat and GO RIGHT!
Decrease temperature of endothermic reaction
Reaction shifts to left
(make more reactants)
Remove a reactant
181kJ + 2HgO ↔ 2 Hg + O2
If we decrease the temperature on this system it is as if we are removing a reactant so it will shift towards the heat and GO LEFT!

41. Changing Temperature—Exothermic
Reaction shifts to left
(get rid of extra products and make more reactants)
Increase temperature of exothermic reaction
Increasing a product
Reaction shifts to right
(make more products)
Decrease temperature of exothermic reaction
Remove a product

42. Some changes have no effect!
Adding a pure solid or liquid reactant or product
They’re not in the equilibrium constant expression
Increasing pressure by adding an inert gas
Changing the volume of a reaction with an equal number of moles of gas on each side of the reaction
The system won’t gain anything by shifting since both sides will cause the same pressure
Adding a catalyst
A catalyst will speed up how fast equilibrium is established—but not the number of reactants and products once it’s at equilibrium

43. Which way will the reaction shift for each of the following changes:
Examples
Example:
Which way will the reaction shift for each of the following changes:
NH4Cl (s)  NH3 (g) + HCl (g)
Removing some NH4Cl
Adding HCl
Adding Ne (g)
Decreasing volume

44. Which way will the reaction shift for each of the following changes:
Examples
Example:
Which way will the reaction shift for each of the following changes:
NH4Cl (s)  NH3 (g) + HCl (g)
Removing some NH4Cl
Adding HCl
Adding Ne (g)
Decreasing volume
No change (it’s a solid)
(Adding a product)
No change (it’s an inert gas)
(Goes to side with least gas moles)

45. Which way will the reaction shift for each of the following changes:
Let’s Practice
Which way will the reaction shift for each of the following changes:
2 SO2 (g) + O2 (g)  2 SO3 (g) an exothermic reaction
Increasing volume
Raising temperature
Adding O2
Removing SO2

46. Which way will the reaction shift for each of the following changes:
Let’s Practice
Which way will the reaction shift for each of the following changes:
2 SO2 (g) + O2 (g)  2 SO3 (g) an exothermic reaction
Increasing volume
Raising temperature
Adding O2
Removing SO2
(Goes to side with most gas moles)
(Energy is a product)
(Adding a reactant)
(Removing a reactant)