1. 3 Chemical Equilibrium COURSE NAME: CHEMISTRY 101
Chapter 3
Chemical
Equilibrium
COURSE NAME: CHEMISTRY 101
COURSE CODE:
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2. Equilibrium is a state in which there are no observable changes as time goes by.
Chemical equilibrium is achieved when:
the rates of the forward and reverse reactions are equal and
the concentrations of the reactants and products remain constant
Physical equilibrium
H2O (l)
H2O (g)
Chemical equilibrium
NO2
N2O4 (g)
2NO2 (g) 1

3. 2 N2O4 (g) 2NO2 (g) equilibrium equilibrium equilibrium Start with NO2
Start with NO2 & N2O4 2

4. 3 N2O4 (g) 2NO2 (g) K = [NO2]2 [N2O4] = 4.63 x 10-3 aA + bB cC + dD
Law of Mass Action 3

5. 4 Equilibrium Will aA + bB cC + dD K = [C]c[D]d [A]a[B]b K > 1
Lie to the left
Favor products
K < 1
Lie to the right
Favor reactants 4

6. Homogenous equilibrium applies to reactions in which all reacting species are in the same phase.
N2O4 (g) NO2 (g)
Kp =
NO2
P 2
N2O4 P
Kc =
[NO2]2
[N2O4]
In most cases
Kc  Kp
aA (g) + bB (g) cC (g) + dD (g)
Kp = Kc(RT)Dn
Dn = moles of gaseous products – moles of gaseous reactants
= (c + d) – (a + b) 5

7. Homogeneous Equilibrium
CH3COOH (aq) + H2O (l) CH3COO- (aq) + H3O+ (aq)
[CH3COO-][H3O+]
[CH3COOH][H2O]
Kc = ′
[H2O] = constant
[CH3COO-][H3O+]
[CH3COOH] =
Kc [H2O] ′
Kc =
General practice not to include units for the equilibrium constant. 6

8. The equilibrium concentrations for the reaction between carbon monoxide and molecular chlorine to form COCl2 (g) at 740C are [CO] = M, [Cl2] = M, and [COCl2] = 0.14 M. Calculate the equilibrium constants Kc and Kp.
CO (g) + Cl2 (g) COCl2 (g)
[COCl2]
[CO][Cl2] =
0.14
0.012 x 0.054
Kc =
= 220
Kp = Kc(RT)Dn
Dn = 1 – 2 = -1
R =
T = = 347 K
Kp = 220 x ( x 347)-1 = 7.7 7

9. 8 The equilibrium constant Kp for the reaction
is 158 at 1000K. What is the equilibrium pressure of O2 if the PNO = atm and PNO = atm? 2
2NO2 (g) NO (g) + O2 (g)
Kp = 2
PNO PO
PNO PO 2
= Kp
PNO PO 2
= 158 x (0.400)2/(0.270)2
= 347 atm 8

10. Heterogenous equilibrium applies to reactions in which reactants and products are in different phases.
CaCO3 (s) CaO (s) + CO2 (g)
[CaO][CO2]
[CaCO3]
Kc = ′
[CaCO3] = constant
[CaO] = constant
[CaCO3]
[CaO]
Kc x ′
Kc = [CO2] =
Kp = PCO 2
The concentration of solids and pure liquids are not included in the expression for the equilibrium constant.
PCO 2
= Kp
PCO 2
does not depend on the amount of CaCO3 or CaO 9

11. 10 Consider the following equilibrium at 295 K:
The partial pressure of each gas is atm. Calculate Kp
and Kc for the reaction?
NH4HS (s) NH3 (g) + H2S (g)
Kp = P
NH3
H2S P
= x =
Kp = Kc(RT)Dn
Kc = Kp(RT)-Dn
Dn = 2 – 0 = 2
T = 295 K
Kc = x ( x 295)-2 = 1.20 x 10-4 10

12. ′ ′ ′ ′ ′′ ′ 11 A + B C + D Kc Kc = [C][D] [A][B] Kc = [E][F] [C][D]
C + D E + F
A + B E + F Kc
[E][F]
[A][B]
Kc =
Kc = Kc ′′ ′ x
If a reaction can be expressed as the sum of two or more reactions, the equilibrium constant for the overall reaction is given by the product of the equilibrium constants of the individual reactions.
N2O4 (g) NO2 (g)
2NO2 (g) N2O4 (g)
= 4.63 x 10-3
K =
[NO2]2
[N2O4]
K =
[N2O4]
[NO2]2 ′ = 1 K
= 216
When the equation for a reversible reaction is written in the opposite direction, the equilibrium constant becomes the reciprocal of the original equilibrium constant. 11

13. Writing Equilibrium Constant Expressions
The concentrations of the reacting species in the condensed phase are expressed in M. In the gaseous phase, the concentrations can be expressed in M or in atm.
The concentrations of pure solids, pure liquids and solvents do not appear in the equilibrium constant expressions.
The equilibrium constant is a dimensionless quantity.
In quoting a value for the equilibrium constant, you must specify the balanced equation and the temperature.
If a reaction can be expressed as a sum of two or more reactions, the equilibrium constant for the overall reaction is given by the product of the equilibrium constants of the individual reactions. 12

14. Chemical Kinetics and Chemical Equilibrium
ratef = kf [A][B]2
A + 2B AB2 kf kr
rater = kr [AB2]
Equilibrium
ratef = rater
kf [A][B]2 = kr [AB2] kf kr
[AB2]
[A][B]2 =
Kc = 13

15. The reaction quotient (Qc) is calculated by substituting the initial concentrations of the reactants and products into the equilibrium constant (Kc) expression. IF
Qc > Kc system proceeds from right to left to reach equilibrium
Qc = Kc the system is at equilibrium
Qc < Kc system proceeds from left to right to reach equilibrium 14

16. Calculating Equilibrium Concentrations
Express the equilibrium concentrations of all species in terms of the initial concentrations and a single unknown x, which represents the change in concentration.
Write the equilibrium constant expression in terms of the equilibrium concentrations. Knowing the value of the equilibrium constant, solve for x.
Having solved for x, calculate the equilibrium concentrations of all species. 15

17. ICE 16 Br2 (g) 2Br (g) Let x be the change in concentration of Br2
At 1280oC the equilibrium constant (Kc) for the reaction
Is 1.1 x If the initial concentrations are [Br2] = M and [Br] = M, calculate the concentrations of these species at equilibrium.
Br2 (g) Br (g)
Let x be the change in concentration of Br2
Br2 (g) Br (g)
Initial (M)
0.063
0.012
ICE
Change (M) -x
+2x
Equilibrium (M) x x
[Br]2
[Br2]
Kc =
Kc =
( x)2 x
= 1.1 x 10-3
Solve for x 16

18. Kc =
( x)2 x
= 1.1 x 10-3
4x x = – x
4x x = 0
-b ± b2 – 4ac  2a
x =
ax2 + bx + c =0
x =
x =
Br2 (g) Br (g)
Initial (M)
Change (M)
Equilibrium (M)
0.063
0.012 -x
+2x x x
At equilibrium, [Br] = x = M
or M
At equilibrium, [Br2] = – x = M 17

19. Le Chatelier’s Principle
If a stress is applied to a system at equilibrium, the system will change to relieve that stress and re –establish equilibrium. It is like the “undo” button on your computer!
Factors that Affect Equilibrium
Concentration
Temperature
Pressure For gaseous systems only!
The presence of a catalyst 18

20. Concentration Changes
Add more reactant  Shift to products
Remove reactants  Shift to reactants
Add more product  Shift to reactants
Remove products  Shift to products
Reaction Quotient
The reaction quotient for an equilibrium system is the same as the equilibrium expression, but the concentrations are NOT at equilibrium!
N2O4(g)  2NO2(g)
Q = [NO2]2 / [N2O4] 19

21. N2O4(g)  2NO2(g) Q = [NO2]2 / [N2O4] 20
Changes in concentration are best understood in terms of what would happen to “Q” if the concentrations were changed.
N2O4(g)  2NO2(g)
Q = [NO2]2 / [N2O4]
Q = Keq at equilibrium
If Q< K then there are too many reactants, the reaction will shift in the forward direction (the products)
If Q>K then there are too many products, the reaction will shift to the reactants. 20

22. Temperature Changes A + B = AB + Heat A + B + heat = AB 21
Consider heat as a product in exothermic reactions
Add heat  Shift to reactants
Remove heat  Shift to products
A + B = AB + Heat
Consider heat as a reactant in endothermic reactions
Add heat  Shift to products
Remove heat  Shift to reactants
A + B + heat = AB 21

23. Pressure Changes N2(g) + 3H2(g) = 2NH3(g) 22
Only affects equilibrium systems with unequal moles of gaseous reactants and products.
N2(g) + 3H2(g) = 2NH3(g)
Increase Pressure
* Stress of pressure is reduced by reducing the number of gas molecules in the container
There are 4 molecules of reactants vs. 2 molecules of products.
Thus, the reaction shifts to the product ammonia. 22

24. PCl5(g) = PCl3(g) + Cl2(g) 23 Decrease Pressure
Stress of decreased pressure is reduced by increasing the number of gas molecules in the container.
There are two product gas molecules vs. one reactant gas molecule.
Thus, the reaction shifts to the products. 23

25. Presence of a Catalyst
A Catalyst lowers the activation energy and increases the reaction rate.
It will lower the forward and reverse reaction rates,
Therefore, a catalyst has NO EFFECT on a system at equilibrium!
It just gets you to equilibrium faster! 24

26. Presence of an Inert Substance
An inert substance is a substance that is not- reactive with any species in the equilibrium system.
These will not affect the equilibrium system.
If the substance does react with a species at equilibrium, then there will be a shift! 25

27. Le Châtelier’s Principle - Summary
Change
Shift Equilibrium
Change Equilibrium
Constant
Concentration
yes no
Pressure
yes* no
Volume
yes* no
Temperature
yes
yes
Catalyst no no
*Dependent on relative moles of gaseous reactants and products 26

28. Question 1
The equilibrium constant Keq for a certain reaction will change if __________ changes.
A) pressure
B) time
C) volume
D) temperature
E) reactant concentrations 27

29. Question 2
What is the relationship of the equilibrium constants for the following two reactions?
(1) 2 NO2(g) ↔ N2O4(g);
(2) N2O4(g) ↔ 2 NO2(g)
A) Keq of reaction (1) is the reciprocal of Keq of reaction (2).
B) Keq of reaction (2) is the reciprocal of Keq of reaction (1)
C) Keq of reaction (1) = Keq of reaction (2)
D) Answers A and B are both correct
E) There is no relationship between the Keqs of these reactions 28

30. Question 3
What is the correct equilibrium constant expression for the following reaction? 2 Cu(s) + O2(g) → 2 CuO(s)
A) Keq = 1/[O2]2
B) Keq = [CuO]2/[Cu]2
C) Keq = [CuO]2/[Cu]2[O2]
D) Keq = [O2]
E) Keq = 1/[O2] 29

31. Question 4
Consider the reaction: 2 SO2(g) + O2(g) ↔ 2 SO3(g). If, at equilibrium at a certain temperature, [SO2] = 1.50 M, [O2] = M, and [SO3] = 1.25 M, what is the value of the equilibrium constant?
A) 5.79
B) 6.94
C) 8.68
D) 0.14
E) None of the above 30

32. Question 5
PCl5 dissociates according to the reaction: PCl5(g) ↔ PCl3(g) + Cl2(g). One mole of PCl5 was placed in one liter of solution. When equilibrium was established, 0.5 mole of PCl5 remained in the mixture. What is the equilibrium constant for this reaction?
A) 0.25
B) 0.5
C) 1.0
D) 2.5
E) None of the above 31

33. Question 6
At elevated temperatures, solid silicon reacts with chlorine gas to form gaseous SiCl4. At some temperature, the equilibrium constant for this reaction is If the reaction is started with 0.10 mol of SiCl4 in a one-liter flask, how much Cl2 will be present when equilibrium is established?
A) 0.18 mol
B) mol
C) mol
D) 0.30 mol 32

34. Question 7
Consider the reaction: CO2(g) + H2(g) ↔ CO(g) + H2O(g), for which Kc = 0.64 at 900 K. If the initial CO2 and of H2 are each M, what will be the equilibrium concentrations of each species after the reaction reaches equilibrium?
A) [CO2] = M; [H2] = 0.044M; [CO] = M; [H2O] = M
B) [CO2] = M; [H2] = M; [CO] = M; [H2O] = M
C) [CO2] = M; [H2] = M; [CO] = M; [H2O] = M
D) [CO2] = M; [H2] = M; [CO] = M; [H2O] = M
E) None of the above 33

35. Question 8
At some temperature, the reaction: 3 ClO- ↔ ClO Cl- has an equilibrium constant Kc = 3.2 x 103. If the components of this reaction are mixed such that their initial concentrations are [Cl-] = 0.05 M; [ClO3-] = 0.32; and [ClO-] = 0.74, is the mixture at equilibrium, yes or no? If the mixture is not at equilibrium in which direction, left to right or right to left, will reaction occur so that the mixture can reach equilibrium?
A) There is not enough information given to answer this question
B) Yes, the mixture is at equilibrium now
C) No. Left to right
D) No. Right to left 34

36. Question 9
Consider the following endothermic reaction: H2(g) + I2(g) ↔ 2 HI(g). If the temperature is increased,
A) more HI will be produced
B) some HI will decompose, forming H2 and I2
C) the magnitude of the equilibrium constant will decrease
D) the pressure in the container will increase
E) the pressure in the container will decrease 35

37. Question 10
Consider the following reaction at equilibrium: NO2(g) + CO(g) ↔ NO(g) + CO2(g). Suppose the volume of the system is decreased at constant temperature, what change will this cause in the system?
A) A shift to produce more NO
B) A shift to produce more CO
C) A shift to produce more NO2
D) No shift will occur 36

38. Question 11
Which of the following statements is incorrect regarding equilibrium?
A) Chemical equilibrium is a reversible process with no net change in concentrations of the products and reactants.
B) Physical equilibrium can not exist between phases.
C) A chemical equilibrium with all reactants and products in the same phase is homogeneous.
D) none of the above 37

39. Question 12
Which of these four factors can change the value of the equilibrium?
A) catalyst
B) pressure
C) concentration
D) temperature
Question 13
Which general rule helps predict the shift in direction of an equilibrium reaction?
A) Le Chatelier's principle
B) Haber process
C) Equilibrium constant
D) Bosch theory
Question 14
There are guideline to help write equilibrium constants.
A) True
B) False 38

40. 39 Question 1 A chemical reaction is at equilibrium when
A) the concentrations of reactants is equal to the concentrations of products
B) the limiting reagent has been completely depleted
C) the rate of the forward reaction equals the rate of the reverse reaction
D) A and B.
 Question 2
What is the correct equilibrium constant expression for the following reaction? 2 NO2(g) ↔ 2 NO(g) + O2(g)
A) Keq = [O2]2[NO]/[NO2]2
B) Keq = [O2][NO]2/[NO2]2
C) Keq = [NO2]2/[NO]2 [O2]
D) Keq = [NO2]2/[NO][O2]2
E) None of the above
 Question 3
The equilibrium constant for the reaction 2 NH3(g) → N2(g) + 3 H2(g) is 3 x 10-3 at some temperature. What is Keq for the reaction 0.5 N2(g) H2(g) ↔ NH3(g) at the same temperature?
A) 0.003
B) 0.05
C) 18
D) 20   39

41. Question 4
Consider the reaction that describes the Haber process for the production of ammonia (NH3): N2(g) + 3 H2(g) ↔ 2 NH3(g) for which Kc at 300 oC is 9.5. Calculate Kp for this reaction at 300 oC.
A) Kp = 4.3 x 10-3
B) Kp = 9.5
C) Kp = 2.1 x 104
D) Kp = 1.6 x 10-2
 Question 5
For the reaction H2(g) + I2(g) ↔ 2 HI(g), Kc = 12.3 at some temperature T. If [H2] = [I2] = [HI] = 3.21 x 10-3 M at that temperature, which one of the following statements is true?
A) The concentration of HI will rise as the system approaches equilibrium
B) The system is at equilibrium, so the concentrations will not change
C) The concentrations of H2 and I2 will increase as the system approaches equilibrium
D) The concentrations of H2 and HI will decrease as the system approaches equilibrium
E) Not enough information is given to answer the question 40

42. Question 6
Kc for the reaction 2 NH3(g) ↔ N2(g) + 3 H2(g) is 3 x 10-3 at some temperature. A 1.0 L mixture containing 1.0 mol of NH3, 0.50 mol of N2, and 0.15 mol of H2 is prepared at this temperature When equilibrium is reached,
A) there will be more N2 and H2 present
B) there will be more NH3 present
C) there will be less N2 but more H2 present.
D) there will be less NH3 and less N2 present
E) No shift will occur   
Question 7
If the equilibrium constant for the reaction PCl5 ↔ PCl3 + Cl2 is 1.0, how many moles of PCl5 must be placed into one liter of solution in order to obtain 0.50 mol of PCl3 when the system reaches equilibrium?
A) 0.25
B) 0.50
C) 0.75
D) 1.00 41

43. Question 8
When the reaction CH3Cl + OH- ↔ CH3OH + Cl- is started with 0.1 mol of CH3Cl and 0.2 mol of OH-, 0.03 mol of CH3OH is present when the system reaches equilibrium. Calculate the equilibrium constant for the reaction.
A) 0.18
B) 0.08 C)
D) 0.30   
Question 9
At some temperature, Kc for the reaction PCl5(g) ↔ PCl3(g) + Cl2(g) is If 0.10 mol of PCl5 and 0.20 mol of PCl3 are added to a 1-L flask, what will be the Cl2 concentration when equilibrium is reached?
A) M
B) 8.7 x 10-3 M
C) M
D) 0.12 M
Question 10
Consider the following reaction: H2(g) + I2(g) ↔ 2 HI(g) for which, at some temperature, Kc = 4.0. When the reaction is started with equimolar quantities of H2 and I2 and equilibrium is reached, 0.20 mol of HI is present. How much H2 was used to start the reaction?
A) 0.10 mol B) 0.23 mol
C) 0.20 mol D) 4.0 mol 42