Ch 12 Equilibrium

General Equilibrium (Kailley/Baverstock 2005) 1

Basic Equilibrium Concepts 1. In the reaction 3 C2H2(g) ↔ C6H6(g) will the ratio of [C2H2] to [C6H6] at equilibrium be 3:1?

2. List conditions that must be true at equilibrium. 3. Ozone (O3) and oxygen gas O2 can exist in equilibrium.

2 O3(g) ↔ 3 O2(g).

If 2 mol of O3 react for every 2 mol of O2 reacting does equilibrium exist in the container?

4. Consider the following reaction:

2 NOCl(g) ↔ 2 NO(g) + Cl2(g).

In a closed system, it is found that 2 moles of NOCl react for every 3 moles of product that react. Is the mixture at equilibrium? Explain.


X(g) ↔ Y(g).

If you start with pure X in the container, and the reaction comes to equilibrium, show a sketch of rates (forward/reverse) vs. time.

6. Consider the following equilibrium reaction:

2 NH3(g) ↔ N2(g) + 3 H2(g)

Why does the [N2] remain unchanged even though the forward reaction continues to occur and produce N2?

7. Consider the following reaction

H2(g) + I2(g) ↔ 2 HI(g)

If you began with HI in a container only, sketch the RATE vs. time graph as equilibrium is reached.


H2(g) + I2(g) ↔ 2 HI(g)

You initially had 3.0 M of H2 and I2 and no HI in the container. The equilibrium concentration of H2 and I2 was found to be 1.4 M. Sketch the concentration vs. time graph for this reaction.


H2(g) + Br2(g) ↔ 2 HBr(g)

You initially had no H2 and Br2 and 3.0 M HBr in the container. The equilibrium [H2] was found to be 1.2 M. Sketch the concentration vs. time graph for this reaction.



H2(g) + F2(g) ↔ 2 HF(g)

You initially had 4.0 MH2 and 2.0 M F2 and no HF in the container. As equilibrium was established, a total of 1.6 M of H2 was reacted (consumed). Sketch the concentration vs. time graph for this reaction.

11. Consider the following concentration vs. time graph sketch for a reaction involving chemicals A, B,

and C. The amount reacted for each chemical is indicated by the variable x.

2x 3x x

a. What were the reactants? Explain. b. Write the overall reaction with the proper COEFFICIENTS (think)

c. How can you tell on a concentration vs. time graph when equilibrium is established?

d. How can you tell on a rate vs. time graph when equilibrium is established?

[ ] M

time

A

B C


Entropy and Enthalpy

NOTE: all reactions are shown as single arrows to prevent you from guessing the answers by looking at the arrows. Many of the reactions will actually be double arrow reactions!

1. In each of the following pairs of substances, select the one which has the greatest entropy, Explain

your answer.

a. H2O(l) or H2O(g) b. Cl2 (g) or 2 Cl-(aq) c. NH3 (l) or NH3 (aq) d. CH3COOH (aq) or CH3COO-

(aq) + H+(aq)

2. In each of the following, decide:

i. Which side is favoured by the tendency to minimum enthalpy. ii. Which side is favoured by the tendency to maximum entropy.

iii. Whether the reaction will be:

A spontaneous reaction which goes to completion or

A non-spontaneous reaction in which NO products are formed or

A spontaneous equilibrium reaction

a. b. c. d.

e. H2SO4 (l) → H2SO4 (aq) + 150 kJ

f. C2H6 (g) → C2H2 (g) + 2 H2 (g) ∆H = + 311 kJ

g. C2H2(g) + Ca(OH)2(aq) → CaC2(s) + 2 H2O(l) ∆H = + 183 kJ

h. 2 C(s) + O2 (g) → 2 CO(g) ∆H = - 221 kJ

PE

(kJ)

Reaction Progress

A(g) + B(g)

5C(g) + 6D(g)

PE

(kJ)

Reaction Progress

4Q(g) + 8R(g)

M(g) + N(g)

PE

(kJ)

Reaction Progress

4X(g) + 7Y(g)

P(g) + Q(g)

PE

(kJ)

Reaction Progress

G(g) + H(g)

5E(g) + 6F(g)


3. Which of the following choices show the changes in entropy and enthalpy as the products are made?

i. Entropy increases enthalpy decreases

ii. Entropy decreases enthalpy increases

iii. Entropy increases enthalpy increases

iv. Entropy decreases enthalpy decreases

a. 2NO2(g) → N2O4(g) H= - 50kJ

b. CaCO3(s) + heat → CaO(s) + CO2(g)

c. N2(g) + 3H2(g) → 2NH3(g) + energy

d. 2NOBr(g) → 2NO(g) + Br2(g) H= + 20kJ 4. Which direction do minimum enthalpy and maximum entropy favour?

a. Liquid nitroglycerine explodes, forming an expanding cloud of gases. b. Solid AgBr is almost insoluble in water; that is, very little Ag+ and Br- are formed when AgBr is

mixed with water.

c. Water and alcohol mix completely in any proportions; that is, they are miscible.

d. The reaction: 3 N2 (g) + Pb (s) → Pb(N3)2 (s) does not occur.

e. When N2O4 (g) is put in a container, some of it decomposes into 2 NO2 (g).

f. Smoke, carbon dioxide and water vapour will not react to make wood and oxygen.


Le Chatelier/Rate Theory: Concentration

1. Create concentration vs. time graphs for the following reactions. Assume the reactions are starting

at equilibrium.

a. 4 HCl(g) + O2(g) ↔ 2 H2O(g) + 2 Cl2(g) + 111.4 kJ At t1 , the [O2] was increased.

b. N2(g) + 3 H2(g) ↔ 2 NH3(g) + 92 kJ At t1 the [H2] was decreased

c. 2 NOCl(g) + 30 kJ ↔ 2 NO(g) + Cl2(g) At t1 , the [Cl2] was decreased.

d. 2 SO2(g) + O2(g) ↔ 2 SO3(g) + 20 kJ At t1 , the [SO3] was increased.

2.

a. Consider the following equation,

Fe3+(aq) + SCN-

(aq) ↔ FeSCN2+(aq) + 25 kJ

From the graph below, describe what happened at t1.

b. Consider the following equation:

N2(g) + 3 H2(g) ↔ 2 NH3(g) + 92 kJ

From the graph below, describe what happened at t1.

3. Create a concentration vs. time graph to describe the following situation.

N2(g) + 3 H2(g) ↔ 2 NH3(g) ∆H = - 92 kJ

This reaction is initially at equilibrium with [N2] = 4.0 M, [H2] = 3.5 M and [NH3] = 2.4 M. At a certain time the equilibrium is disturbed by adding 0.6 M NH3. The new equilibrium has a [N2] of 4.2 M.

FeSCN2+

SCN-

Fe3+

[ ] M

time t1

H2

N2

NH3 [ ] M

time t1


4. What (if anything) is incorrect with the concentration vs. time graphs for the following situations? Create the proper graph if necessary.

a. 2 NOCl(g) + 30 kJ ↔ 2 NO(g) + Cl2(g) At t1 , the [NOCl] was decreased.

b. N2(g) + 3 H2(g) ↔ 2 NH3(g) + 92 kJ At t1, the [NH3] was increased.


2 NOCl(g) +30 kJ ↔ 2 NO(g) + Cl2(g) Brown

a. Describe what happens when the [NO] is decreased (using Le Chateliers’s principle).

b. Describe what happens when the [NO] is decreased (using Rate theory).


4 HCl(g) + O2(g) ↔ 2 H2O(g) + 2 Cl2(g) + 111.4kJ

a. Describe what happens when the [HCl] is decreased (using Le Chateliers’s principle).

b. Describe what happens when the [HCl] is decreased (using Rate theory).

7. Create rate vs. time graphs for questions #1 a-d. Show the rates during the initial equilibrium to the new equilibrium that is established. Identify which line represents the forward and which represents the reverse rate.

NOCl

NO

Cl2

[ ] M

time t1

N2

H2

NH3 [ ] M

time t1


Le Chatelier/Rate Theory: Temperature 1. Create concentration vs. time graphs for the following reactions. Assume the reactions are starting

at equilibrium. a. 4 HCl(g) + O2(g) ↔ 2 H2O(g) + 2 Cl2(g) + 111.4 kJ At t1 , the temperature was increased.

b. N2(g) + 3 H2(g) ↔ 2 NH3(g) ∆H = -92 kJ At t1 the temperature was decreased.

c. 2 NOCl(g) + 30 kJ↔ 2 NO(g) + Cl2(g) At t1 , the temperature was decreased.

d. 2 SO2(g) + O2(g) ↔ 2 SO3(g) + 20 kJ At t1 , the temperature was increased.

2. Use the following reactions and graph to answer the questions below.

a. Fe3+(aq) + SCN-

(aq) ↔ FeSCN2+(aq) + 25 kJ What happened at t1?

b. 2 NH3(g) + 92 kJ ↔ N2(g) + 3 H2(g) What happened at t1?

3. Create a concentration vs. time graph to describe the following situation. Explain whether

temperature was increased or decreased.

N2(g) + 3 H2(g) ↔ 2 NH3(g) + heat

This reaction is initially at equilibrium with [N2] = 4.0 M, [H2 ] = 3.5 M and [NH3] = 2.4 M. At a certain time the equilibrium is disturbed by changing the temperature. The new [H2] is 2.9 M.

Fe3+

SCN-

FeSCN2+

[ ]M

time t1

H2

N2

NH3 [ ]M

time t1


4. What (if anything) is incorrect with the concentration vs. time graphs for the following situations. Create the proper graph if necessary.

a. 2 NOCl(g) + 30 kJ ↔ 2 NO(g) + Cl2(g) At t1 , the temperature was decreased. b. N2(g) + 3 H2(g) ↔ 2 NH3(g) ∆H = -92 kJ At t1, the temperature was increased.


2 NOCl(g) +30 kJ ↔ 2 NO(g) + Cl2(g) Brown

a. Describe what happens when the temperature is decreased (using Le Chatelier’s principle). b. Describe what happens when the temperature is decreased (using Rate theory).


4 HCl(g) + O2(g) ↔ 2 H2O(g) + 2 Cl2(g) + 111.4 kJ

a. Describe what happens when the temperature is increased (using Le Chatelier’s principle). b. Describe what happens when the temperature is increased (using Rate theory).

7. Create rate vs time graphs for questions #1 a-d. Show the rates during the initial equilibrium to the

new equilibrium that is established. Identify which line represents the forward and which represents the reverse rate.

8. Consider the following reaction: A + B ↔ C + D

yellow blue.

When the temperature is decreased, the solution goes from yellow to blue. Explain if the forward reaction is endothermic or exothermic, using Le Chatelier’s Principle

9. Consider the following reaction: A + B ↔ C + D red green

When the temperature is increased, the solution goes from red to green. Explain if the forward reaction is endothermic or exothermic, using Le Chatelier’s principle.

Cl2

NO

NOCl [ ] M

time t1

H2

N2

NH3 [ ] M

time t1


Le Chatelier/Rate Theory: Volume/Catalysts 1. Create concentration vs. time graphs for each of the following reactions. Assume the reactions are

starting at equilibrium.

a. 4 HCl(g) + O2(g) ↔ 2 H2O(g) + 2 Cl2(g) + 111.4 kJ At t1 , the volume was increased.

b. 110 kJ + N2(g) + 3 H2(g) ↔ 2 NH3(g) At t1 the volume was decreased

c. 2 NOCl(g) + 30 kJ↔ 2 NO(g) + Cl2(g) At t1 , the volume was decreased.

d. H2(g) + Br2(g) ↔ 2 HBr(g) + 20 kJ At t1 , the volume was increased.

e. Fe3+(aq) + SCN-

(aq) ↔ FeSCN2+(aq) + 25 kJ At t1 , the container volume was increased.

2. Use the following reactions and graphs to answer the questions.

a. 2 NOCl(g) + 30 kJ↔ 2 NO(g) + Cl2(g) What happened at t1? b. N2(g) + 3 H2(g) ↔ 2 NH3(g) + heat What happened at t1?

3. What (if anything) is incorrect with the concentration vs. time graphs for the following situations?

Create the proper graph if necessary.

a. 2 NOCl(g) + 30 kJ↔ 2 NO(g) + Cl2(g) At t1 , the volume was decreased.

NOCl

NO

Cl2

[ ] M

time t1

NH3

H2

N2 [ ] M

time t1

NOCl

NO

Cl2 [ ] M

time t1


b. N2(g) + 3 H2(g) ↔ 2 NH3(g) + 92 kJ At t1, the volume was decreased.


2 NOCl(g) + 30 kJ↔ 2 NO(g) + Cl2(g) Brown

a. Describe what happens when the volume is decreased (using Le Chateliers’s principle). b. Describe what happens when the volume is decreased (using Rate theory).


4 HCl(g) + O2(g) ↔ 2 H2O(g) + 2 Cl2(g) + 111.4 kJ

a. Describe what happens when the volume is decreased (using Le Chateliers’s principle). b. Describe what happens when the volume is decreased (using Rate theory).

6. Create rate vs. time graphs for questions #1 a-d. Show the rates during the initial equilibrium to the

new equilibrium that is established. Identify which line represents the forward and which represents the reverse rate.

7. Create a rate vs. time graph and a concentration vs. time graph showing what happens when a

catalyst is added to the reaction:

4 HCl(g) + O2(g) ↔ 2 H2O(g) + 2 Cl2(g) + 111.4 kJ.

NH3

H2

N2

[ ] M

time t1


Le Chatelier’s Principle: Summary Problems 1. Use Le Chatelier’s Principle to describe the effect of the following changes on the position of the

equilibrium. (Just describe the stress and shift)

a. The equilibrium is: N2O3(g) ↔ NO(g) + NO2(g)

i. increase the [NO] ii. increase the [N2O3] iii. increase the pressure by decreasing the volume iv. add a catalyst

b. The equilibrium is : 2 H2(g) + 2 NO(g) ↔ N2(g) + 2H2O(g)

i. decrease the [N2] ii. decrease the [NO] iii. increasing the volume

c. The equilibrium is: 2 CO(g) + O2(g) ↔ 2 CO2(g) + 566 kJ.

i. increase the temperature ii. increase the [O2] iii. increase the pressure by decreasing the volume

d. The equilibrium is: I2(g) + Cl2(g) → 2 ICl(g); H = + 35.0 kJ

i. decrease the temperature ii. decrease the [Cl2] iii. decreasing the volume

For each exercise 2-4, describe the effect on the concentration of the bold substance by the following changes. Write ↑ for increase, ↓ for decrease, or nc for no change.

2. N2(g) + 3 H2(g) ↔ 2 NH3(g) H = - 92 kJ

a. increase the [N2] b. increase the temperature c. increase the volume d. add a catalyst

3. 2 HF(g) ↔ F2(g) + H2(g) H = + 536 kJ

a. decrease the temperature b. decrease the [H2] c. decrease the volume d. increase the partial pressure of H2

4. SnO2(s) + 2 CO(g) ↔ Sn(s) + 2 CO2(g) H = + 13 kJ

a. increase the temperature b. add a catalyst c. increase the [CO] d. add Kr at a constant volume e. add SnO2


5. Construct concentration vs. time graphs for each of the following reactions and stresses. Note: In Exercises I-III the relative positioning of the molecules is not relevant; simply place them on the graph so the reactants are separated from the products. The only thing required here is to show how a substance’s concentration will change.

a. H2(g) + I2(g) ↔ 2 HI(g) + 52 kJ

i. increase the temperature ii. inject some H2(g) iii. decrease the volume iv. add a catalyst

b. 2 SO2(g) + O2(g) ↔ 2 SO3(g) H = - 197 kJ

i. inject some SO2 ii. increase the volume iii. decrease the temperature iv. increase the [SO3]

c. CO(g) + H2O(g) ↔ CO2(g) + H2(g) H = - 41 kJ

i. inject some CO2(g) ii. remove some of the H2O(g) with a very rapidly acting drying agent iii. increase the temperature iv. decrease the pressure by increasing the volume

From the graphs, determine what change must have been imposed on the equilibrium. 6. PCl5(g) + 92.5 kJ ↔ PCl3(g) + Cl2(g)

a. b.

7. H2O(g) + Cl2O(g) ↔ 2 HOCl(g) + 70 kJ.

a. b.

PCl5

PCl3

Cl2 [ ] M

time t1

PCl5

PCl3

Cl2 [ ] M

time t1

HOCl

Cl2O

H2O [ ] M

time t1

HOCl

Cl2O

H2O [ ] M

time t1


Disturbing Equilibrium

You can print a copy of this lab sheet on the website. Purpose: To observe and explain the colour change that occurs when different stresses are applied to a reaction. Procedure: Part I: Indicator Equilibrium 1. Add 5 drops of bromthymol blue and 5 mL of water to a 50 mL beaker. Record the colour. 2. Add 5 drops of 0.1 M HCl to the beaker, swirl the beaker then record the colour change. 3. Add 20 drops of 0.1 M NaOH to the same beaker, swirl the beaker then record the colour change. 4. Dispose of the solutions down the sink. Part II: Thiocyanatoiron(III) ion Equilibrium 1. Add 1 drop of 0.2 M FeCl3, 1 drop of 0.2 M KSCN and 10 mL of water to a 50 mL beaker. Record

the colour.

2. Use a medicine dropper to add 5 drops of the above solution to a spotwell. To the same spotwell add 4 drops of 0.2 M Fe(NO3)3 and record the colour change.

3. To the same spotwell add 4 drops of 1.0 M NaOH and carefully record all observations.

4. Dispose of the solutions down the sink. Data and Observations: Table I: HIn ↔ H+ + In- Table II: Fe3+ + SCN- ↔ FeSCN2+ Yellow Blue Yellow Dark Red

Observations

Initial Colour

0.1 M HCl added

0.1 M NaOH added

Observations

Initial Colour

0.2 M Fe(NO3)3 added

1.0 M NaOH added


Equilibrium Expression Keq

1. Write the equilibrium expressions (Keq) for the following.

a. 2 ICl(g) ↔ I2(g) + Cl2(g)

b. N2(g) + O2(g) ↔ 2 NO(g)

c. 3 O2(g) ↔2 O3(g)

d. 2 Bi3+(aq) + 3 H2S(g) ↔ Bi2S3(s) + 6 H+

(aq)

e. CaCO3(s) ↔ CaO(s) + CO2(g)

f. CaC2(s) + 2 H2O(l) ↔ C2H2(g) + Ca(OH)2(s)

g. C6H6(l) + Br2(l) ↔ C6H5Br(l) + HBr(g)

h. Cu(s) + 2 Ag+(aq) ↔ Cu2+

(aq) + 2 Ag(s)

i. 4 NH3(g) + 5 O2(g) ↔ 6 H2O(g) + 4 NO(g)

j. H2(g) + ½ O2(g) ↔ H2O(l)

2. Write the Keq expression for:

a. N2O4(g) ↔ 2NO2(g)

b. 2NO2(g) ↔ N2O4(g)

c. Examine the relationship between the Keq expressions for equations (a) and (b) of this question. If Keq = 10.0 for equation a then what would be the value for Keq for equation (b)?

3. Write the Keq expression for:

a. SO2(g) + ½ O2(g) ↔ SO3(g)

b. 2 SO2(g) + O2(g) ↔ 2SO3(g)

c. Examine the relationship which exists between he Keq expressions for equations (a) and (b) of this question. If Keq = 3.0 for equation (a), what would be the value of Keq for equation (b)?

4. Consider the reaction: CaCO3(s) + CO2(g) + H2O(l) ↔ Ca2+

(aq) + 2 HCO3-(aq) + 40 kJ

Which way will the equilibrium shift if:

a. more CO2(g) is added? b. CaCO3(s) is powdered? c. Ca2+

(aq) is removed? d. heat is added?

5. Rearrange the following equations to solve in terms of the concentrations indicated in bold.

a. Keq = [H3O+][F-]/[HF] d. Keq = [NO2]

2/[NO]2[O2] g. Keq = [NH3]2/[N2][H2]

3 b. Keq = [H3O

+][F-]/[HF] e. Keq = [NH3]2/[N2][H2]

3 h. Keq = [PCl3]4/[P4][Cl2]

6 c. Keq = [NO2]

2/[NO]2[O2] f. Keq = [N2O4]/[NO2]2


6. Consider the following equilibrium reactions.

i. 2 NO2(g) ↔ N2O4(g) Keq = 2.2 ii. Cu2+

(aq) + 2 Ag(s) ↔ Cu(s) + 2 Ag+(aq) Keq = 1 x 10-15

iii. Pb2+(aq) + 2 Cl-(aq) ↔ PbCl2(s) Keq = 6.3 x 104

iv. SO2(g) + 2 ½ O2(g) ↔ SO3(g) Keq = 110

a. Which equilibrium favours products to the greatest extent?

b. Which equilibrium favours reactants to the greatest extent? 7. In the reaction A(g) + B(g) ↔ C(g) + D(g) + 100 kJ, what happens to the value of Keq if the

temperature is increased? Explain using Le Chatelier’s Principle. 8. If the value of Keq decreases when the temperature decreases, is the forward reaction exothermic

or endothermic? Explain using Le Chatelier’s Principle. 9. In the reaction P(g) + Q(g) + 150 kJ ↔ Y(g) + Z(g), what happens to the value of Keq if the [Y] is

increased? 10. If the value of Keq increases when the temperature decreases is the reaction exothermic or

endothermic. Explain using Le Chatelier’s Principle. 11. In the equilibrium reaction : AgCl(s) + 17 kJ ↔ Ag+

(aq) + Cl-(aq), which way will the equilibrium shift and what is the effect on the value of Keq when:

a. AgNO3 is added? b. the temperature is decreased? c. more AgCl(s) is added?

12. A(aq) + 2 B(g) ↔ 2 C(aq) + 2 D(aq) has a Keq value of 0.25 at 100oC and a Keq value of 0.15 at 200oC.

Is the forward reaction endothermic or exothermic? Explain using Le Chatelier’s Principle. 13. Examine the following graphs for the equilibrium 3 O2 (g) ↔ 2 O3 (g) [ ] (M) Time T = 250 oC [ ] (M) Time Is the forward reaction endothermic or exothermic, as written? Use the graphs and Le Chatelier’s Princple to explain.

O2

O3

T = 50 oC

O2

O3


Equilibrium Calculations Set 1

Calculating Keq


CO(g) + H2O(g) ↔ CO2(g) + H2(g)

At equilibrium at a certain temperature, it is found that the [CO] =1.02 M, [H2O] =1.50 M, [CO2] =2.25 M, and [H2] =1.75 M. Calculate the value of Keq.


N2(g) + 3 H2(g) ↔ 2 NH3(g)

At equilibrium at a certain temperature, it is found that the [N2] = 0.0345 M, [H2] = 0.0225 M, and [NH3] = 3.22 M. Calculate the value of Keq?

3. A 1.0 L reaction vessel contains 0.750 mol of CO and 0.275 mol of H2O. After one hour,

equilibrium is reached according to the equation CO(g) + H2O(g) ↔ CO2(g) + H2(g). Analysis shows 0.250 mol of CO2 present at equilibrium. What is Keq for the reaction?

4. A 5.0 L reaction vessel was initially filled with 6.0 mol of SO2, 2.5 mol of NO2 and 1.0 mol of SO3.

After equilibrium was established according to the equation SO2(g) + NO2(g) ↔ SO3(g) + NO(g), the vessel was found to contain 3.0 mol of SO3. What is Keq for the reaction?

5. Consider the equilibrium N2(g) + 3 H2(g) ↔2 NH3(g)

a. At a certain temperature 3.0 mol of N2 and 2.0 mol of H2 are put into a 5.0 L container. At equilibrium the concentration of NH3 is 0.020 M. Calculate Keq for the reaction.

b. At a different temperature, 6.0 mol of NH3 were introduced into a 10.0 L container. At equilibrium 2.0 mol of NH3 were left. Calculate Keq for the reaction.

6. When 0.50 mol of NOCl was put into a 1.0 L flask and allowed to come to equilibrium, 0.10 mol of

Cl2 was found. What is Keq for the reaction: 2 NOCl(g) ↔ 2 NO(g) + Cl2(g)?

Calculating Equilibrium Concentrations

1. Keq = 5.0 at a certain temperature for the reaction 2 SO2(g) + O2(g) ↔ 2 SO3(g). A certain amount of SO3 was placed in a 2.0 L reaction vessel. At equilibrium the vessel contained 0.30 mol of O2. What concentration of SO3 was originally placed in the vessel?

2. Keq = 35.0 for the reaction PCl5(g) ↔ PCl3(g) + Cl2(g). At equilibrium, [PCl5] = 1.34 x 10-3 M and

[PCl3] = 0.205 M in a certain vessel, what is the equilibrium concentration of Cl2? 3. A student obtained the following data at 25oC while studying the equilibrium.

2 Tl+(aq) + Cd(s) ↔ 2 Tl(s) + Cd2+(aq)

Volume Moles Tl+ Moles Cd2+

1.00 L 0.316 0.414

5.00 L ? 0.339

Calculate the number of moles of Tl+ present in the second data set.


4. Keq = 7.5 for 2 H2(g) + S2(g) ↔ 2 H2S(g). A certain amount of H2S was added to a 2.0 L flask and allowed to come to equilibrium. At equilibrium, 0.072 mol of H2 was found. How many moles of H2S were originally added to the flask?

5. Keq = 49.5 for H2(g) + I2(g) ↔ 2 HI(g) at a certain temperature. If 0.250 mol of H2 and 0.250 mol of I2

are placed in a 10.0 L vessel and permitted to react, what will be the concentration of each substance at equilibrium?

6. The equilibrium constant for the reaction N2(g) + 3 H2(g) ↔ 2 NH3(g) is 3.0 at a certain temperature.

Enough NH3 was added to a 5.0 L container such that at equilibrium the container was found to contain 2.5 mol of N2. How many moles of NH3 were put into the container?

7. Keq = 1.00 for N2O2(g) + H2(g) ↔ N2O(g) + H2O(g). If 0.150 mole of N2O and 0.250 mol of H2O were

introduced into a 1.00 L bulb and allowed to come to equilibrium, what concentration of N2O2 was present at equilibrium?

EIRE Calculations

1. A reaction mixture at equilibrium, CO(g) + H2O(g) ↔ CO2(g) + H2(g), contains 1.00 mol of H2,

2.00 mol of CO2, 2.00 mol of CO and 2.00 mol of H2O in a 2.00 L bulb. If 1.00 mol of H2 is added to the system, calculate the [CO] which will exist when equilibrium is regained.

2. A reaction mixture at equilibrium, CO2(g) + H2(g) ↔ CO(g) + H2O(g), contained 4.00 mol of CO2, 1.50 mol of H2, 3.00 mol of CO and 2.50 mol of H2O in a 5.0 L container. How many mole of CO2 would have to be removed from the system in order to reduce the amount of CO to 2.50 mol?

3. A reaction mixture at equilibrium, H2(g) + I2(g) ↔ 2 HI(g), contains 0.150 mol of H2, 0.150 mole of I2 and 0.870 mol of HI in a 10.0 L vessel. If 0.400 mol of HI is added to this system and the system is allowed to come to equilibrium again, what will be the new concentrations of H2, I2, and HI?

4. A reaction mixture, 2 NO(g) + O2(g) ↔ 2 NO2(g), contained 0.240 mol of NO, 0.0860 mol of O2 and 1.20 mol of NO2 when at equilibrium in a 2.00 L bulb. How many moles of O2 had to be added to the mixture to increase the number of moles of NO2 to 1.28 when equilibrium was re-established?

5. A reaction mixture, 2 ICl(g) + H2(g) ↔ I2(g) + 2 HCl(g), was found to contain 0.500 mol of ICl, 0.0560 mol of H2, 1.360 mol of I2 and 0.800 mol of HCl at equilibrium in a 1.00 L bulb. How many moles of ICl would have to be removed in order to reduce the [HCl] to 0.680 M when equilibrium is re-established?

6. (Nasty!) Keq = 100 at a certain temperature for CH4(g) + 2 H2S(g) ↔ CS2(g) + 4 H2(g). Some CH4 and H2S were introduced into a 1.0 L bulb and at equilibrium 0.10 mol of CH4 and 0.30 mol of H2S were found. What was [CS2] at equilibrium?


Q calculations

1. Consider the following reaction

CH4(g) + 2 H2S(g) ↔ CS2(g) + 4 H2(g) Keq= 2.5 x 103.

A 10.0 L reaction vessel contains 2.0 mol CH4, 3.0 mol CS2, 3.0 mol H2 and 4.0 mol H2S. Is the mixture at equilibrium, and if not, which direction will it shift in order to establish equilibrium.


H2O(g) + CH4(g) ↔ CO(g) + 3 H2(g) Keq= 4.7

The [H2O] = 0.035 M, [CH4] = 0.050 M, [CO] = 0.15 M, [H2] = 0.20 M.In which direction does the reaction proceed to establish equilibrium?


2 NO(g) + O2(g) ↔ 2 NO2(g) Keq= 6.9 x 105

A 5.0 L vessel is filled with 0.060 mol NO, 1.0 mol O2 and 0.80 mol NO2. Is the reaction mixture at equilibrium, and if not which direction will it proceed to get to equilibrium? Will the pressure of the vessel increase or decrease?


SO2(g) + NO2(g) ↔ NO(g) + SO3(g) Keq= 85.0

Initially 0.100 mol SO2, 0.100 mol NO2, 0.0800 mol NO, and 0.0800 mol SO3 are placed in a 10.0 L container. What will the equilibrium concentration of all species be?


NO2(g) + NO(g) ↔ N2O(g)+ O2(g) Keq= 0.914

A mixture was prepared initially by adding 0.200 mol of each gas into a 5.00 L container. What will the equilibrium concentrations be? What changes will occur in pressure as equilibrium is established?


Equilibrium Calculations Set 2

Calculating Keq 1. For the reaction:

2 SO2(g) + O2(g) ↔ 2 SO3 (g)

10.0 mol of SO3 is added to a 2.50 L container and at equilibrium [SO2] is 0.50 M. Calculate Keq. 2. For the reaction:

N2(g) + 3 H2(g) ↔ 2 NH3(g)

5.00 M of NH3 and 6.00 M of N2 are added to a 5.00 L container. At equilibrium 1.50 M of H2 exists. Calculate Keq.

3. For the reaction:

2 C(s) + O2(g) ↔ 2 CO(g)

21.0 g of carbon is mixed with 52.0 g of O2 in a 3.00 L container. At equilibrium [CO] is 0.402 M.

a. Calculate Keq. b. Calculate the final mass of carbon.


Fe3+(aq) + SCN-

(aq) ↔ FeSCN2+(aq)

Initially, 50.0 mL of 0.20 M Fe3+ is added to 30.0 mL 0.20 M SCN-. At equilibrium, the concentration of FeSCN2+ is found to be 0.050 M. Calculate Keq. (Careful this involves a dilution.)

Calculating Equilibrium Concentrations


CO(g) + H2O(g) CO2(g) + H2(g) Keq = 4.06

1.00 M of CO and H2O are placed in a 10.0 L flask and establish equilibrium. Calculate the mass of CO that exists at equilibrium.


Fe3+ + SCN- FeSCN2+ Keq = 12.0

4.836 g of Fe(NO3)3 and 2.916 g of KSCN are added to 200 mL of water. Calculate the equilibrium

concentrations. Watch the significant figures in the quadratic equation (,,,) break up the steps.


EIRE Problems

1. PCl5(g) ↔ PCl3(g) + Cl2(g)

At equilibrium this reaction is found to contain [PCl3] = 1.25 M, [Cl2] = 0.750 M and [PCl5] = 2.26 M. How many moles of PCl5 must be removed from a 10.0 L container to decrease the [Cl2] to 0.525 M?

2. 2 SO2Cl(g) ↔ 2 SO2(g) + Cl2(g) Keq = 2.55

At equilibrium this reaction contains [SO2] = 0.215 M and [SO2Cl] = 0.0817 M. What mass of Cl2 must be added to a 6.00 L container to increase the [SO2Cl] to 0.103 M?

3. SbCl3(g) + Cl2(g) ↔ SbCl5(g)

At equilibrium an 8.50 L flask contains 9.05 g Cl2, 23.69 g SbCl3 and 140.0 g SbCl5. What concentration of SbCl3 must be removed to decrease the concentration of SbCl5 to 0.0515 M?

4. 2 ICl(g) + H2(g) ↔ I2(g) + 2 HCl(g)

The above reaction was found to contain 1.00 mol of ICl, 0.1120 mol H2, 2.720 mol I2 and 1.60 mol HCl at equilibrium in a 2.000 L container. How many moles of ICl would have to be removed in order to reduce the concentration of HCl to 0.680 M when equilibrium is re-established?

Q Problems

1. N2O2(g) + H2(g) N2O(g) + H2O (g) Keq = 1.00

If 0.150 M N2O2, 0.150 M H2, 0.250 M N2O and 0.250 M H2O are introduced into a 1.00 L flask and allowed to establish equilibrium then calculate the equilibrium concentrations of N2O2 and H2O.

2. Br2(g) 2 Br(g) Keq = 0.0011

Initially 6.3 x 10-2 M Br2 and 1.2 x 10-2 M Br are added to a container. Is this reaction at equilibrium? If not explain in what direction this reaction must shift to establish equilibrium.

. .

3. CO(g) + H2O(g) CO2(g) + H2(g) Keq = 10.0

Initially 8.00 mol of CO, 8.00 mol H2O, 5.00 mol CO2 and 5.00 mol H2 are added to a 2.00 L flask. Calculate the equilibrium concentrations.

4. S(s) + O2(g) SO2 (g) Keq = 0.0234

3.20 g of O2, 12.0 g S and 9.62 g of SO2 are added to a 500 mL container. Calculate the equilibrium concentrations.

5. 2 NO(g) N2(g) + O2(g) Keq = 20.00

0.2000 M N2, 0.4000 M O2 and 0.5000 M NO are added to a container. Calculate the equilibrium concentrations.


Solutions: Definitions, Review and Applications of Equilibrium Definitions: 1. Define the following:

a. saturated solution

b. unsaturated solution

c. supersaturated solution

d. solute

e. solvent

f. solubility

g. electrolyte

h. non-electrolyte

i. ionic solution

j. molecular solution

2. Use the above definitions to classify the following solutions as:

a. a strong electrolyte, weak electrolyte or non-electrolyte b. an ionic or molecular solution

i. NaCl

ii. BaCO3

iii. HNO3

iv. Pb(NO3)2

v. CH3COOH

vi. CH3CH2OH

Review: 3. What substances make up an ionic compound? 4. What substances make up a molecular compound? 5. Write dissociation reactions for the following solutions:

a. K2SO4

b. CuCl2

c. Fe3(PO4)2

d. PbS

e. (NH4)2CO3

f. Cr(OH)3

g. Au(NO3)3

h. Sr(OH)2

6. Consider the dissociation equations from question 5.

a. For which of the above reactions can we simply use the mole ratio to calculate the ion concentrations? Explain.

b. How do you think we will calculate the concentration for the other reactions?

7. Calculate the concentration of each ion.

a. 5.42 M Fe2(SO4)3 b. 10.0 g of K2CO3 is added to 500 mL of water

c. 0.500 mol of Mg(OH)2 is dissolved in 1.00 L of water


8. Write balanced, full ionic (complete) and net ionic reactions for the following:

a. 1.0 M calcium sulphide reacts with 1.0 M sodium hydroxide.

b. A solution of ammonium phosphate reacts with a solution of barium chloride

c. Solid lead (II) nitrate is dissolved in water and mixed with a solution of strontium bromide.

d. Solutions of strontium hydroxide and iron (III) sulphate are mixed together. 9. Outline an experiment to separate a mixture of:

a. Ba2+, Cu2+ and Pb2+

b. S2-, SO42- and CO3

2- 10. A mixture may contain: Al3+, Ag+, Sr2+ or Mg2+. Experimentally how would you determine

what is present? 11. A mixture may contain one or more of the ions: S2-, OH-, Cl-, CO3

2-. Experimentally how would you determine what is present using only the reagents: AgNO3, Ba(NO3)2, Cu(NO3)2 and Sr(NO3)2?

Applications of Equilibrium: 12. Consider the equilibrium:

Fe3+ + 3 OH- ↔ Fe(OH)3(s) + heat

a. Explain why this reaction is able to establish equilibrium.

b. When heat is added, what happens to the mass of solid? Explain using Le Chatelier's Principle.

c. When HNO3 added to the above equilibrium, what happens to the solubility of Fe(OH)3?

Explain using Le Chatelier's Principle.

d. When Fe(NO3)3 is added, what is the direction of equilibrium shift? Explain using rate theory.


Solubility Problems Set 1

Calculating Solubility

1. Calculate the solubility of lead (II) iodate.

2. What mass of calcium fluoride dissolves in 250 mL of water? Ksp = 1.7 x 10-10 3. Calculate the concentration of silver ions in a 500 mL solution of Ag2CrO4. 4. 0.500 g of silver carbonate is added to 800 mL of water. What mass of solid remains

undissolved? 5. Calculate the solubility of strontium fluoride in g/mL. 6. A saturated solution of aluminum hydroxide has a [Al3+] = 0.00250 M calculate the

concentration of hydroxide ions. Ksp = 3.7 x 10-15 7. What is the [OH-] and [Zn2+] ions in a saturated solution of zinc hydroxide? Ksp= 5.3 x 10-15 8. What [S2-] is required to just start precipitation of copper (II) sulphide from a 0.20 M solution of

copper (II) chloride? 9. What is the maximum [F-] that can be added to a 3.0 x 10-3 M solution of calcium nitrate before

a calcium fluoride precipitate forms? Ksp =1.5 x 10-10

Calculating Ksp 10. The solubility of barium sulphate is 0.0091 g/L. Calculate Ksp.

11. A 200 ml saturated solution of silver acetate is evaporated. After all the water is removed

1.60 g of solid remains. Calculate Ksp. 12. An experiment is performed to find the solubility product constant of an unknown compound

with the general formula X2Y3. 2000 mL of a saturated solution containing X2Y3 is evaporated to remove all the water. The following data is collected:

Solubilty data for X2Y3

Volume of saturated solution 2000 mL

Mass of beaker 285.63 g

Mass of beaker and X2Y3 after 1st heating 292.75 g

Mass of beaker and X2Y3 after 2nd heating 292.64 g

Mass of beaker and X2Y3 after 3rd heating 292.64 g

Molecular mass of X2Y3 124.2 g/mol

Use this data to calculate Ksp for the ionic compound X2Y3

13. A student determined that 0.0981 g of lead (II) fluoride was dissolved in 200 mL of saturated

lead (II) fluoride solution. What is the Ksp for lead (II) fluoride?


TIP Problems 14. 2.00 L of 0.015 M calcium nitrate is mixed with 3.00 L of 0.045 M sodium oxalate. Will a

precipitate form?

15. A solution contains [Bi3+] = 9.85 x 10-15 M and [S2-] = 1.48 x 10-14 M. If the Ksp value is 2.93 x 10-70 for bismuth (III) sulphide then does a precipitate form?

16. 0.0681 g of strontium hydroxide and 0.0431 g of sodium fluoride are dissolved in 500 mL of

water. Will a precipitate form? 17. 50 mL of 5.10 x 10-4 M strontium hydroxide is mixed with 200 mL of 1.23 x 10-2 M

magnesium nitrate. Will a precipitate form? 18. Equal volumes of 0.0216 M lead (II) nitrate and 0.0412 M sodium bromide are mixed together.

Will a precipitate form?

Common Ion Effect Problems 19. What is the solubility of silver chloride (in g/L) in 6.5 x 10-3 M silver nitrate solution?

20. Calculate the number of grams of zinc sulphide that will dissolve in 300 mL of 0.050 M

zinc nitrate. 21. Calculate the solubility of cobalt (II) hydroxide (Ksp = 2.5 x 10-16) in 0.0032 M solution of

strontium hydroxide. 22. Calculate the solubility of lead (II) chloride in 0.020 M aluminum chloride. 23. How many grams of sodium fluoride must be added to 1.00 L of solution to reduce the

solubility of barium fluoride to 6.8 x 10-4 M? Ksp = 1.7 x 10-6 24. A 500 mL portion of 0.0020 M sodium oxalate solution is able to dissolve 0.47 g of

magnesium oxalate. What is the Ksp for magnesium oxalate?


Solubility Demo

A copy of this demo sheet can be found on the website. Purpose: To observe the ways in which solubility can be changed. Procedure: Teacher Demo Data and Observations: Data Table I: CuCO3 Solubility Changes.

Species Added Stress Ion Observations Direction of Shift

Original -------------- ----------------------

Add Na2CO3

Add HNO3

Add KI

Add NH3

Add H2O

Questions: Explain in full sentences. Include an equation at the top of every answer or graph that involves Le Chatelier’s Principle or Rate Theory. Include the Ksp expression when explaining changes to the Ksp value. 1. State the three ways to remove ions from solution? 2. What is the common ion effect? How does it affect solubility? 3. Will adding water change the solubility of CuCO3? Explain. 4. Use Le Chatelier’s Principle to explain how the solubility of CuCO3 changes when HNO3 is

added. 5. Graph the concentration changes that occur when:

a. at t1 add HNO3 b. at t2 add NH3 c. at t3 add K2CO3 d. at t4 add heat CuCO3

Concentration (M) Cu2+ CO3

2-

Time(s)

6. How will the Ksp value change when heat is added? Explain using the graph. 7. How will the Ksp value change when K2CO3 is added? Explain using the graph. 8. Use Rate Theory to explain how the solubility changes when NH3 is added. 9. Graph the rate changes that occur when: (label the axes)

a. at t1 add HNO3 b. at t2 add NH3 c. at t3 add K2CO3 d. at t4 add heat


Mixed Solubility Problems 1. Write the Ksp expression for the following:

a. PuO2CO3(s) ↔ PuO22+ + CO3

2-

b. Ni3(PO4)2 (s) ↔ 3 Ni2+ + 2 PO43-

2. Write the equilibrium reactions for each of the following:

a. Ksp = [Fe3+]2[SO42-]3

b. Ksp = [Hg22+][I-]2

3. Calculate the solubility:

a. lead (II) bromide

b. magnesium arsenate Ksp = 2.1 x 10-20

4. Calculate the solubility of magnesium hydroxide in:

a. Pure water

b. 0.015 M magnesium chloride

c. 0.217 M potassium hydroxide

5. Predict whether a precipitate will form:

a. 0.015 mol of magnesium nitrate is mixed with 0.0072 mol sodium carbonate in 1.00 L water.

b. Equal volumes of 0.0076 M silver nitrate and 0.021 M sodium sulphate. Ksp = 1.4 x 10-4

c. 10 mL of 0.205 M chromium nitrate is mixed with 40 mL of 7.75 x 10-12 M sodium hydroxide. Ksp = 6.3 x 10-31

6. A solution of potassium iodide is slowly added to a solution that 0.10 M in both Pb2+ and Ag+.

a. Which precipitate should form first?

b. What [I-] is required for the second cation to begin precipitation?

c. What concentration of the first cation remains in solution when the second cation just begins to precipitate?

7. Should a precipitate of magnesium fluoride (Ksp = 3.7 x 10-8) occur if a 17.5 mg sample of

magnesium chloride hexahydrate is dissolved in 325 mL of 0.045 M potassium fluoride solution?

8. Outline an experiment to separate:

a. lead (II) ions and aluminum ions

b. copper (II) ions and magnesium ions

c. chloride ions and sulphate ions


Equilibrium and Solubility Review Set 1

1. Explain the term dynamic equilibrium.

2. For the reaction: CO(g) + 2 H2(g) ↔ CH3OH(g) ∆ H = - 234 kJ

a. Graph

i. Addition of H2 at t1 ii. Removal of CH3OH at t2 iii. Addition of heat at t3

b. What happens to the overall [CO] when:

i. Heat is removed ii. CO is added iii. H2 is added iv. Volume is increased

3. Are the following reactions capable of establishing equilibrium? Explain

a. PCl5(g) → PCl3(g) + Cl2(g) ∆H = (+)

b. Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g) ∆ H = (-) 4. Consider the equilibrium: A + B ↔ C + D

Yellow Blue

When heat is added to this system it turns blue. Is this reaction exothermic or endothermic? Explain using Le Chatelier’s Principle.

5. A 3.00 L reaction flask was initially filled with 5.00 mol of CO2 and 6.00 mol of H2 and allowed to establish equilibrium. At equilibrium the [H2] is 1.50 M, what is the value of Keq?

CO(g) + H2O(g) ↔ CO2 (g) + H2(g)

6. 1.50 mol of CO2 and 7.50 mol of C are added to a 20.0 L container, at equilibrium the

concentration of CO is 0.0700 M. Calculate Keq.

C (s) + CO2(g) ↔ 2 CO(g)

7. 0.600 mol of X and 0.400 mol of Y are reacted in a 2.00 L container yielding an equilibrium concentration of 0.080 M Z. Calculate Keq.

X (g) + 2Y (g) ↔ 2 Z(g)

8. 6.90 M of HCl is initially added to a closed system. What is the concentration of H2 at

equilibrium?

H2 (g) + Cl2 (g) ↔ 2 HCl (g) Keq = 0.018


9. At 800 oC the equilibrium constant for the following reaction is 0.279. At a different temperature the equilibrium constant is 0.100. Is this different temperature higher or lower than 800oC? Explain using Le Chatelier’s Principle.

CO2 (g) + H2 (g) ↔ CO(g) + H2O(g) ∆ H = + 42.6 kJ

10. Initially concentrations of H2, Br2 and HBr are 1.50 M, 1.50 M and 0.321 M respectively. What

are the equilibrium concentrations?

H2 (g) + Br2(g) ↔ 2 HBr (g) Keq = 1.50

11. Calculate the Ksp value of Cr(OH)3 if it has a solubility of 1.30 x 10-6 g/L. 12. 355 mL of 1.3 x 10-2 M aluminum nitrate is mixed with 265 mL of 3.2 x 10-4 M

strontium hydroxide. Will a precipitate form if Ksp = 1.9 x 10-33 for aluminum hydroxide?

13. Use Le Chatelier’s Principle to explain why the solubility of Cu3(PO4)2 decreases in 0.20 M Na3PO4.

14. Calculate the solubility of Cu3(PO4)2 in 0.20 M Na3PO4. Ksp = 1.4 x 10-40

15. How does the solubility and Ksp value of barium carbonate change when nitric acid is added?


Equilibrium and Solubility Review Set 2

1. Consider the following equilibrium

CO(g) + H2O(g) ↔ CO2 (g) + H2(g) Keq = 4.0

The equilibrium concentrations in a 2.0 L container are: [CO] = 0.10 M [CO2] = 0.20 M [H2O] = 0.10 M [H2] = 0.20 M

How many moles of CO2 must be injected into the flask in order to increase the concentration of CO to 0.20 M?


2 NO(g) + O2(g) ↔ 2 NO2(g)

In an experiment, 0.30 mol of NO and 0.80 mol of NO2 are placed in a 5.0 L flask at 10oC. When equilibrium is reached, it is found that the [O2] is 0.020 M. Calculate the value of Keq.


2 NO(g) + Cl2(g) ↔ 2 NOCl(g)

A mixture consisting of 2.00 moles of NO and 2.00 moles of Cl2 are placed in a container. When the system reaches equilibrium, 30 % of the original NO is reacted. Calculate Keq.


A2(g) + B2(g) ↔ 2 AB(g)

There are found to be 2.00 mol of A2, 2.00 mol of B2 and 4.00 mol of AB in a 10.0 L flask. If 1.00 mol of AB is added to the system then what is the new equilibrium concentration of AB.

5. Consider the equilibrium:

PCl3(g) + Cl2(g) ↔ PCl5(g) ∆H = - 88 kJ

a. What direction does entropy and enthalpy favour? Explain.

b. How will the [Cl2] be affected by?

i. adding PCl3

ii. adding PCl5

iii. raising temperature

iv. decreasing the volume of the flask

v. adding a catalyst

c. How will the above changes affect the value of Keq?


6. The following is a list of Ksp values for various silver salts:

Silver bromate, 5.77 x 10-5 Silver carbonate, 6.15 x 10-12

Silver dichromate, 2.0 x 10-7 Silver thiocyanate, 1.16 x 10-12

Silver hydroxide, 1.52 x 10-8 Silver bromide, 7.7 x 10-13

Silver chloride, 1.56 x 10-10 Silver iodide, 1.5 x 10-16

Silver chromate, 9.0 x 10-12

a. Which of the salts will give the highest concentration of silver ions in a saturated solution? b. Calculate this concentration.

7. What is the solubility of calcium fluoride in g/L. Ksp = 4.9 x 10-11 8. At a certain temperature the Ksp for magnesium hydroxide is 1.8 x 10-11.

a. What is the solubility of magnesium hydroxide? b. What is the concentration of hydroxide ions in solution?

9. A given sample of hard water has a [Ca2+] of 1.0 x 10-3 M.

a. What is hard water? b. If the Ksp of calcium fluoride is 1.7 x 10-10, what is the maximum [F-] that can be attained in

this sample of water before calcium fluoride would precipitate? 10. Will a precipitate form when equal volumes of 2.0 x 10-3 M aluminum chloride are mixed with

4.0 x 10-2 M sodium hydroxide? Ksp for the precipitate is 3.7 x 10-15 11. Outline a scheme to separate a solution of Ba2+, Ca2+ and Fe3+. 12. What is the solubility of magnesium hydroxide in 0.10 M sodium hydroxide if the Ksp is

7.1 x 10-12? 13. Use Le Chatelier’s Principle to explain whether adding the following ions to a saturated

solution of silver ethanoate will increase or decrease the solubility?

a. adding sodium ethanoate

b. adding potassium iodide

c. adding ammonia

d. adding potassium nitrate

e. adding nitric acid

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Ch 12 Equilibrium