1. Kinetics applies to the speed of a reaction, the concentration of product that appears (or of reactant that disappears) per unit time.
Equilibrium applies to the extent of a reaction, the concentration of product that has appeared after an unlimited time, or once no further change occurs.
At equilibrium: rateforward = ratereverse
A system at equilibrium is dynamic on the molecular level; no further net change is observed because changes in one direction are balanced by changes in the other. 1

2. CHEMICAL EQUILIBRIUM - RATES OF REACTION kF Reactants  products kB
Chemical reactions are a dynamic process, that is, reactions involve both forward and reverse processes.
Chemical Equilibrium is reached by a reaction mixture
when the rates of forward and reverse reactions becomes equal.
kF = kB
NO net change appears obvious although the system is
still in constant motion. 2

3. Reaction Dynamics Because the reactant concentration
decreases, the forward reaction slows.
As the products accumulate, the
reverse reaction speeds up.
As the forward reaction proceeds
it makes products and uses reactants.
Once equilibrium is established,
the forward and reverse reactions
proceed at the same rate, so the
concentrations of all materials
stay constant.
Initially, only the forward
reaction takes place.
Eventually, the reaction proceeds
in the reverse direction as fast as
it proceeds in the forward direction.
At this time equilibrium is established.
Rate
Rate Forward
Rate Reverse
Time 3

4. Reaching equilibrium on the macroscopic and molecular levels.
N2O4(g) NO2(g) 4

5. H2(g) + I2(g)  2 HI(g)
at time 0, there are only reactants in the mixture, so only the forward reaction can take place
[H2] = 8, [I2] = 8, [HI] = 0
at time 16, there are both reactants and products in the mixture, so both the forward reaction and reverse reaction can take place
[H2] = 6, [I2] = 6, [HI] = 4 5 5

6. H2(g) + I2(g)  2 HI(g)
at time 32, there are now more products than reactants in the mixture − the forward reaction has slowed down as the reactants run out, and the reverse reaction accelerated
[H2] = 4, [I2] = 4, [HI] = 8
at time 48, the amounts of products and reactants in the mixture haven’t changed – the forward and reverse reactions are proceeding at the same rate – it has reached equilibrium 6

7. K  equilibrium constant aA + bB  cC + dD K = Products Reactants
LAW OF MASS ACTION
K  equilibrium constant
aA + bB  cC + dD
K = Products
Reactants
K = [C]c [D]d = Q reaction quotient
[B]b [A]a
Write the equilibrium equation for:
a. H2O + H2O  H3O+ + OH-
b. 4NH3(g) (g)  2N2(g) + 6H2O(g)
c. N2(g) + 3H2(g)  2NH3(g) 7

8. HOMOGENEOUS EQUILIBRIA
Workshop Question E1
HOMOGENEOUS EQUILIBRIA
H2(g) + I2(g)  2HI(g)
Kc = then Kp = _________
2O3(g)  3O2(g)
Kp = Kc = ________
HETEROGENEOUS EQUILIBRIA
2H2O(l)  H3O+(aq)_ + OH-(aq)
C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(g) 8

9. HETEROGENEOUS EQUILIBRIA Substance in more than one phase
1. CaCO3(s)  CaO(s) + CO2(g)
Kc =
How do the [ ] of solid express?
Pure solids & liquid have constant [conc. ] ? 9

10. solids do not change their concentrations
The reaction quotient for a heterogeneous system.
solids do not change their concentrations 10

11. c) solid CaCO3 & PCO2 > Kp d) CaCO3 & CaO
Q. Each of the mixtures listed below was placed in a closed container and allowed to stand. Which of these mixtures is capable of attaining the equilibrium, expressed by 1
a) pure CaCO3
b) CaO & PCO2 > Kp
c) solid CaCO3 & PCO2 > Kp
d) CaCO3 & CaO 11

12. LE CHATELIER’S PRINICIPLE
“If a system at equilibrium is disturbed by a change in
temperature, pressure, or concentration of one of its
components, that system will shift it’s equilibrium position as to ‘counteract’ the effect of the disturbance.”
Equilibrium can be disturbed by:
- adding or removing components
- a change in pressure
- a change in volume
- a change in temperature 12

13. PREDICTING THE DIRECTION OF THE SHIFT I. CATALYST
A catalyst increases the rate at which equilibrium is achieved but not the composition of the
equilibrium mixture.
II. THE REACTION QUOTIENT
Q < K: the reaction shifts to the products
Q > K: the reaction shifts to the reactants
III. CHANGES IN VOLUME
Reducing the volume of a gas at equilibrium causes
the system to shift in the direction that reduces
the number of moles of gas. 13

14. The effect of pressure (volume) on an equilibrium system.+
lower P
(higher V)
more moles of gas
higher P
(lower V)
fewer moles of gas 14

15. The Effect of Volume Changes on Equilibrium
Since there are more gas molecules on the reactants side of the reaction, when the pressure is increased the position of equilibrium shifts toward the products.
When the pressure is decreased by increasing the volume, the position of equilibrium shifts toward the side with the greater number of molecules – the reactant side. 15 15

16. (a) CaCO3(s) CaO(s) + CO2(g)
Sample Problem
Predicting the Effect of a Change in Volume (Pressure) on the Equilibrium Position
PROBLEM:
How would you change the volume of each of the following reactions to increase the yield of the products.
(a) CaCO3(s) CaO(s) + CO2(g)
(b) S(s) + 3F2(g) SF6(g)
(c) Cl2(g) + I2(g) ICl(g)
PLAN:
When gases are present a change in volume will affect the concentration of the gas. If the volume decreases (pressure increases), the reaction will shift to fewer moles of gas and vice versa.
SOLUTION: 16

17. IV. CHANGES IN TEMPERATURE
When heat is added to a system, the equilibrium
shifts in the direction that absorbs heat. Cooling has the opposite effect and shifts the equilibrium towards the side which produces heat.
Endothermic Reactions:
heat + Reactants  Products
An increase in temperature leads to a shift towards the products, a decrease leads to a shift towards the reactants. (Keq increases)
Exothermic Reactions:
Reactants  Products + heat
An increase in temperature leads to a shift towards reactants. (Keq decreases) 17

18. The Effect of Temperature Changes on Equilibrium18

19. (a) CaO(s) + H2O(l) Ca(OH)2(aq) H0 = -82kJ
Sample Problem
Predicting the Effect of a Change in Temperature on the Equilibrium Position
PROBLEM:
How does an increase in temperature affect the concentration of the underlined substance and Kc for the following reactions?
(a) CaO(s) + H2O(l) Ca(OH)2(aq) H0 = -82kJ
(b) CaCO3(s) CaO(s) + CO2(g) H0 = 178kJ
(c) SO2(g) S(s) + O2(g) H0 = 297kJ
SOLUTION: 19

20. Practice - Le Châtelier’s Principle
The reaction 2 SO2(g) + O2(g)  2 SO3(g) with H° = -198 kJ is at equilibrium. How will each of the following changes affect the equilibrium concentrations of each gas once equilibrium is re-established?
adding more O2 to the container
condensing and removing SO3
compressing the gases
cooling the container
doubling the volume of the container
warming the mixture
adding the inert gas helium to the container
adding a catalyst to the mixture 20

21. C) increase the total pressure by adding N2
Workshop question E2a
N2O4(g)  2 NO2(g) H = 58 kJ
In which direction will the equilibrium shift when each of the following changes are made to a system at equilibrium?
A) add N2O4
B) remove NO2
C) increase the total pressure by adding N2
D) increase the volume of the container
E) decrease the temperature 21

22. PCl5(g)  PCl3(g) + Cl2(g) H = 88 kJ
Workshop Question E2b
PCl5(g)  PCl3(g) + Cl2(g) H = 88 kJ
In which direction will the equilibrium shift when each of the following changes are made to a system at equilibrium?
A) add Cl2
B) temperature is increased
C) the volume of the reaction system is decreased
D) PCl5 is added
E) a catalyst is added 22

23. Q - The Reaction Quotient
At any time, t, the system can be sampled to determine the amounts of reactants and products present. A ratio of products to reactants, calculated in the same manner as K tells us whether the system has come to equilibrium (Q = K) or whether the reaction has to proceed further from reactants to products (Q < K) or in the reverse direction from products to reactants (Q > K).
We use the molar concentrations of the substances in the reaction. This is symbolized by using square brackets - [ ].
For a general reaction aA + bB cC + dD where a, b, c, and d are the numerical coefficients in the balanced equation, Q (and K) can be calculated as
Q =
[C]c[D]d
[A]a[B]b 23

24. Reaction direction and the relative sizes of Q and K.
reactants products
reactants products
Equilibrium:
no net change
Reaction Progress
Reaction Progress 24

25. The change in Q during the N2O4-NO2 reaction.25

26. Q, K, and the Direction of Reaction26 26

27. Predicting the direction of reaction: Q > K forms more reactants 
Q = K equilibrium
Q < K forms more products 
Note: 1. Kf = 1 Kr
2. K = Kn If the balanced equation is multiplied
by a factor then the K (& Q) is
multiplied by the exponent. 27

28. The range of equilibrium constants
small K
large K
intermediate K 28

29. UNDERSTANDING THE DIRECTION OF REACTIONS AND Keg
The following reaction is a means of “fixing” nitrogen:
N2(g) + O2(g)  2 NO(g)
A. If the value for Q at 25°C is 1 x 10-30, describe
the feasibility of this reaction for Nitrogen
fixation.
B. Write the equilibrium expression, Kc
C. Write the equilibrium expression for
2NO(g)  N2(g) + O2(g)
D. Determine the Kc for “C” 29

30. Workshop Question E3
1) The equilibrium constant is given for one of the reactions below. Determine the value of the missing equilibrium constant. Which reaction is the most feasible at this temperature and explain why?
  2 SO2(g) + O2(g) ⇌ 2 SO3(g) Kc = 1.7 × 106
SO3(g) ⇌ 1/2 O2(g) + SO2(g) Kc = ?
2) The equilibrium constant is given for two of the reactions below. Determine the value of the missing equilibrium constant.
  A(g) + 2B(g) ⇌ AB2(g) Kc = 59
AB2(g) + B(g) ⇌ AB3(g) Kc = ?
A(g) + 3B(g) ⇌ AB3(g) Kc = 478
3) Consider the following reaction and its equilibrium constant: I2(g) ⇌ 2I(g) Kp = 0.209
 Write the equilibrium expression. If a reaction mixture contains 0.89 atm I2 and 1.77 atm I. Which of the following statements is TRUE concerning this system?
A) The reaction will shift in the direction of reactants.
B) The reaction quotient will increase.
C) The reaction will shift in the direction of products.
D) The equilibrium constant will decrease.
E) The system is at equilibrium.

31. K as either Kc or Kp
Kc = the equilibrium constant using concentrations.
Kp = the equilibrium constant using pressure
Kp = Kc(RT) n gas
n is the difference between the number of moles of products and moles of reactants
R = L atm / mol K 31

32. Kp = Kc(RT) n gas R = 0.082 L atm/ mol K
Kc vs. Kp
Kp = Kc(RT) n gas R = L atm/ mol K
For the following write Kp then Kc for:
N2(g) + 3H2(g)  2NH3(g)
Calculate Kp if Kc = 32

33. Workshop E4 on Kc vs. Kp Write Kp then Kc for: 1. N2O4(g)  2NO2(g)
2. Calculate Kc if Kp = for Q#1
2SO3(g)  2SO2(g) + O2(g)
if Kc= 4.07x10-3, what is Kp? 33

34. Predicting the direction of reaction: Q > K forms more reactants 
Q = K equilibrium
Q < K forms more products 
Note: 1. Kf = 1 Kr
2. K = Kn If the balanced equation is multiplied
by a factor then the K (& Q) is
multiplied by the exponent. 34

35. The range of equilibrium constants
small K
large K
intermediate K 35

36. Calculating Variations on Q and K
[C]c[D]d
Qc =
aA + bB cC + dD
[A]a[B]b
Q’ = 1/Qc
cC + dD aA + bB
Qc’ = (Qc)n n
aA + bB cC + dD
For a sequence of equilibria, Koverall = K1 x K2 x K3 x … 36

37. CALCULATING THE Keq at Equilibrium
1. In one experiment, Haber introduced a mixture of H2 & N2 into a reaction vessel and allowed the system to attain chemical equilibrium at 472°C. The equilibrium mixture of gases were analyzed and found to contain M H 2, M N2, and M NH3. Calculate Keq. ? 38

38. CALCULATING THE Keq at Equilibrium
2. For the Haber process:
N2(g) + 3H2(g)  2NH3(g)
Kp = x at 500°C
a) If an equilibrium mixture of the three gas have partial pressures of atm for H2 and atm for N2, what is the partial pressure of NH3?
b) If an equilibrium mixture of the three gas started with partial pressures of atm for H2 and atm for N2, what is the equilibrium partial pressure of NH3? 39

39. Workshop Eb1 on CALCULATING THE Keq at Equilibrium
1. In another experiment, Haber introduced ammonia into a reaction vessel and allowed the system
to attain chemical equilibrium at 472°C. The
equilibrium mixture of gases were analyzed and was
found to contain 1.56 M H 2, 9.23 M N2, and
M NH3. Calculate Keq.
2. Nitryl Chloride, NO2Cl, is in equilibrium in a closed container with NO2 and Cl2.
2 NO2Cl(g)  2 NO2(g) + Cl2(g)
Calculate Keq if [NO2Cl]eq= M, [NO2]eq= M & [Cl2]eq = M ? 40

40. PREDICTING DIRECTION OF REACTION
1. Predicting the direction of reaction
Q = reaction quotient
at equil Q = K Q > K  species on Rt (prod)
(no net Rx) react to form left
K = [Equil] Q = [Non Equil]
Q < K  forms more products
Goal: Calculate Q to determine state of Rx, equil, more product or more reaction 41

41. I.C.E. & PREDICTING DIRECTION OF REACTION
3a. Sulfur Trioxide decomposes at High temperature in a sealed container.
2 SO3(g)  2 SO2(g) + O2 (g)
Initially the vessel is filled at 1000K with SO3(g) at a concentration of 6.09 x 10-3M. At equilibrium, the [SO3] is 2.44 x 10-3M. Calculate Kc.
3b. Calculate the value for Q and predict the direction in which the reaction will proceed towards equilibrium if the initial concentrations are:
[SO3] = 2.0 x 10-3M
[SO2] = 5.0 x 10-3M
[O2] = 3.0 x 10-2M 42

42. Workshop Eb2 on I.C.E. & Predicting Direction
1. A mixture of 5.0 x 10-3 mol of H2 and 1.0 x 10-2 mol of I2 is placed in a 5.0L container at 448°C and allowed to come to equilibrium. Analysis of this equilibrium mixture shows that the [HI] is 1.87 x 10-3 M.
Calculate Kc: H2(g) + I2(g) ↔ 2HI(g)
2. At 448°C the equilibrium constant Kc for the reaction below is 50.5
H2(g) + I2(g) ↔ 2HI(g)
Predict how the reaction will proceed to reach
equilibrium if the initial amount of HI is 2.0 x 10-2 mol,
H2 is 1.0 x 10-2 mol, and I2 is 3.0 x 10-2 mol in a 2.00 L container. 43

43. Finding Equilibrium Concentrations When Given the Equilibrium Constant and Initial Concentrations or Pressures
first decide which direction the reaction will proceed
compare Q to K
define the changes of all materials in terms of x
use the coefficient from the chemical equation for the coefficient of x
the x change is + for materials on the side the reaction is proceeding toward
the x change is  for materials on the side the reaction is proceeding away from
solve for x
for 2nd order equations, take square roots of both sides or use the quadratic formula
may be able to simplify and approximate answer for very large or small equilibrium constants 44

44. Calculating Keq
Quadratic formula
4. A 3.00 L flask is filled with 1.00 mol of H2 and 2.00 mol I2 at 448°C and Kc is 50.5, what are the equilibrium concentrations of H2, I2 & HI?
Successive Approximation
2B. For the Haber process: N2(g) + 3H2(g)  2NH3(g)
Kp = x at 500°C
If an equilibrium mixture of the three gas started with partial pressures of atm for H2 and atm for N2, what is the partial pressure of NH3 (eq)? 45

45. PCl5(g)  PCl3(g) + Cl2(g)
Calculating Keq
5. The equilibrium constant for the Haber process at 472°C is Kc = A 2.00 L flask is filled with mol of ammonia and is then allowed to reach equilibrium at 472°C. What are the equilibrium concentrations?
N2(g) H2(g)  2 NH3(g)
6. For the reaction
PCl5(g)  PCl3(g) + Cl2(g)
at a certain temperature Kc equals What will happen when 0.10 mol of PCl5, 1.0 mol of PCl3, and l.5 mol of Cl2 are added to a 2.0-L container and the system is brought to the temperature at which Kc=450. What are the equilibrium concentrations? 46

46. Calculating Keq
7. A mixture of mol of hydrogen and 0.01 mol iodine is placed in a 5.0 L container at 448 oC and allowed to reach equilibrium. Analysis of the equilibrium mixture shows the [HI] to be 1.87x10-3M. Calculate K at this temperature.
8. In the decomposition of HI into its elements; what would be the equilibrium partial pressures of each constituent, if the initial pressure of each introduced into the reaction chamber was 1.00 M? Kp = at this temperature. 47

47. Workshop Eb3 on Calculating Keq
1) The decomposition of ammonia is: 2 NH3(g) → N2 (g) + 3 H2 (g). If Kp is 1.5 × 103 at 400°C, what is the partial pressure of ammonia at equilibrium when N2 is atm and H2 is 0.15 atm?
2) At a certain temperature, Kc equals 1.4 × 102 for the reaction: 2 CO(g) + O2 (g) ⇌ 2 CO2 (g).
If a L flask contains mol of CO2 and mol of O2 at equilibrium, how many moles of CO are also present in the flask?
3) At a certain temperature the equilibrium constant, Kc, equals 0.11 for the reaction:
2 ICl(g) ⇌ I2 (g) + Cl2 (g).
What is the equilibrium concentration of ICl if mol of I2 and mol of Cl2 are initially mixed in a 2.0-L flask?
4) Kc is 1.67 × 1020 at 25°C for the formation of iron(III) oxalate complex ion:
Fe3+ (aq) + 3 C2O42-(aq) ⇌ [Fe(C2O4)3]3-(aq).
If M Fe3+ is initially mixed with 1.00 M oxalate ion, what is the concentration of Fe3+ ion at equilibrium? 48

48. EFFECT OF VARIOUS DISTURBANCES ON AN EQUILIBRIUM SYSTEM
DISTURBANCE NET DIRECTION OF REACTION EFFECT ON VALUE OF K
Concentration
Increase (reactant) Toward formation of product None
Decrease (reactant) Toward formation of reactant None
Pressure (volume)
Increase P Toward formation of lower
amount (mol) of gas1 None
Decrease P Toward formation of higher
amount (mol) of gas None
Temperature
Increase T Toward absorption of heat Increases H°rxn>0
Decreases if H°rxn<0
Decrease T Toward release of heat Increases H°rxn<0
Decreases H°rxn>0
Catalyst added None; rates of forward and
reverse reactions increase equally None 49

49. The formation of methanol is an important industrial
VAN’T HOFF EQUATION
Changes in K due to T
In K2 = H°rxn ( )
K R T T2
R = J/mol K
T = Kelvin
The formation of methanol is an important industrial
reaction in the processing of new fuels. At 298K,
Kp = x 104 for the reaction
CO(g) H2(g)  CH3OH(l)
If H°rxn = kJ/mol CH3OH, calculate Kp at 0°C. 50

50. - The van’t Hoff Equation The Effect of T on K ln K2 K1 = - H0rxn R 1
R = universal gas constant
= J/mol*K ln K2 K1
= -
H0rxn R 1 T2 T1 -
K1 is the equilibrium constant at T1
Temperature Dependence ln k2 k1
= - Ea R 1 T2 T1 - ln P2 P1
= -
Hvap R 1 T2 T1 - ln K2 K1
= -
H0rxn R 1 T2 T1 - 51

51. Workshop Eb4 – putting it all together
CH4(g) + CO2(g)  2CO(g) H2(g)
A. What is the theoretical yield of H2 when an equimolar mixture of CH4 and CO2 with a total pressure of 20.0 atm reaches equilibrium at 1200K at which Kp = x 106?
B. What is the [H2]eq for this system at 1300K, at which Kp = x 107?
C. Use the Van’t Hoff equation to find H°rxn. 52