Name:________________________________
_____
Chemistry 30
Equilibrium, Acids and Bases Workbook
Chemical Equilibrium
1. For each of the following, write the chemical reaction equation with the appropriate
equilibrium arrows.
a) pH measurements indicate that acetic acid in vinegar is approximately 1% ionized
into hydrogen ions and acetate ions.
b) Quantitative analysis of the reaction of sodium sulphate and calcium chloride
solutions shows that the products are favoured.
c) Aluminum sulphate solution reacts quantitatively with a sodium hydroxide solution.
d) An analysis of the reaction between sodium hypochlorite and ammonium nitrate
shows that the reactants are favoured.
2. Write the equilibrium law for each of the following chemical reaction equations.
a) 2 SO2(g) + O2(g) ⇌ 2 SO3(g)
b) 2 NO2(g) ⇌ 2 NO(g) + O2(g)
c) N2(g) + 3 H2(g) ⇌ 2 NH3(g)
d) Na2SO4(aq) + CaCl2(aq) ⇌ 2 NaCl(aq) + CaSO4(s)
3. Chlorine and carbon monoxide gases are mixed in a 1.00 L container. Initially, 1.5 mol
of chlorine was present with excess carbon monoxide. The percent reaction is 53%.
a) Write the chemical reaction equation including the appropriate equilibrium arrows
and the percent reaction.
b) Write the equilibrium law for this reaction.
c) At equilibrium, 0.80 mol of COCl2(g), 1.75 mol of carbon monoxide and 0.70 mol of
chlorine were present. Calculate the equilibrium constant.
4. In an experiment at 200C, 0.500 mol/L of hydrogen bromide gas is placed in a sealed
container and it decomposes into hydrogen gas and bromine gas.
a) Write the equilibrium equation and law for this reaction.
b) The equilibrium concentrations in this system are: [HBr(g)] = 0.240 mol/L,
[H2(g)] = [Br2(g)] = 0.130 mol/L. Calculate the equilibrium constant.
Chem 30 Equilibrium, Acids & Bases Workbook
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Graphical Analysis
The Haber-Bosch process of the industrial production of ammonia involves the equilibrium
N2(g) + 3 H2(g)  2 NH3(g) + 92.2 kJ.
In a laboratory experiment designed to study this equilibrium, a chemical engineer injects
0.20 mol of N2(g) and 0.60 mol of H2(g) into a 1.0 L flask at 500C. She records her analysis of
the flask at 5 s intervals in the table shown.
Time (s)
0
5
10
15
20
25
N2(g)
0.20
0.14
0.11
0.10
0.10
0.10
Concentration (mol/L)
H2(g)
0.60
0.42
0.33
0.30
0.30
0.30
NH3(g)
0.00
0.12
0.18
0.20
0.20
0.20
Analyze the data by:
1. Draw a graph of the concentrations of N2(g), H2(g) and NH3(g) versus time on the graph
paper below. Include a legend with your graph.
2. State the time required for equilibrium to be established
3. Calculate the equilibrium constant for this reaction…showing all work.
Chem 30 Equilibrium, Acids & Bases Workbook
2
Le Châtelier’s Principle
1. Nitrogen monoxide, a major air pollutant, is formed in automobile engines from the
endothermic reaction of nitrogen gas and oxygen gas.
a) Write the equilibrium reaction equation including the term “energy” in the equation.
b) What is the direction of the equilibrium shift if the concentration of oxygen is
increased?
c) What is the direction of the equilibrium shift if the pressure is increased?
d) What is the direction of the equilibrium shift if the temperature is decreased?
e) Gasoline burns better at higher temperatures. What is one disadvantage of operating
automobile engines at higher temperatures?
2. In a sealed container, nitrogen monoxide gas and oxygen gas react to form nitrogen
dioxide gas and are allowed to come to equilibrium. The reaction of nitrogen monoxide
and oxygen is exothermic.
a) Write the equilibrium reaction equation including the term “energy” in the equation.
b) What is the direction of the equilibrium shift if the temperature is decreased.
c) What is the direction of the equilibrium shift if the [NO(g)] is decreased.
d) What is the direction of the equilibrium shift if the [NO2(g)] is increased.
e) What is the direction of the equilibrium shift if the volume of the system is decreased.
3. The following is an equilibrium mixture:
Fe3+(aq) + SCN(aq) ⇌ FeSCN2+(aq)
faint yellow
colourless
red
Predict the colour change in the mixture when each of the following changes is made:
a) a crystal of KSCN(s) is added to the system
b) a crystal of FeCl3(s) is added to the system
c) a crystal of NaOH(s) is added to the system
Chem 30 Equilibrium, Acids & Bases Workbook
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Le Châtelier’s Principle
―from Hebden: Chemistry 12
1. Use Le Châtelier’s Principle to describe the effect of the following changes on the
position of the equilibrium:
A. The equilibrium is: N2O3(g) ⇌ NO(g) + NO2(g)
a) increase the [NO]
b) increase the [N2O3]
c) add a catalyst
d) increase the pressure by decreasing the volume
B. The equilibrium is: 2 H2(g) + 2 NO(g) ⇌ N2(g) + 2 H2O(g)
a) decrease the [N2]
b) decrease the pressure by increasing the volume
c) decrease the [NO]
C. The equilibrium is: 2 CO(g) + O2(g) ⇌ 2 CO2(g) + 566 kJ
a) increase the temperature
b) add a catalyst
c) increase the [O2]
D. The equilibrium is: I2(g) + Cl2(g) ⇌ 2 ICl(g) H = +35.0 kJ
a) decrease the temperature
b) decrease the [Cl2]
c) increase the pressure by decreasing the volume
2. Describe the effect (increase, decrease or no change) on the concentration of the boldfaced substance by each of the given changes.
A. The equilibrium is: N2(g) + 3 H2(g) ⇌ 2 NH3(g)
a) increase the [N2]
b) increase the volume
c) increase the temperature
d) add a catalyst
H = 92.0 kJ
B. The equilibrium is: 2 HF(g) ⇌ F2(g) + H2(g)
a) decrease the [H2]
b) decrease the volume
c) decrease the temperature
H = +536.0 kJ
C. The equilibrium is: SnO2(s) + 2 CO(g) ⇌ Sn(s) + 2 CO2(g) H = +13.0 kJ
a) increase the [CO]
b) add a catalyst
c) increase the temperature
Chem 30 Equilibrium, Acids & Bases Workbook
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3. Show the equilibrium adjustment in each of the following situations graphically:
***Note that the relative positioning of the molecules is not relevant
A. The equilibrium is: H2(g) + I2(g) ⇌ 2 HI(g) + 52 kJ
a) increase the temperature
b) decrease the volume
c) inject some H2(g)
d) add a catalyst
B. The equilibrium is: 2 SO2(g) + O2(g) ⇌ 2 SO3(g)
a) inject some SO2(g)
b) decrease the temperature
c) increase the volume
d) increase the [SO3(g)]
H = 197.0 kJ
C. The equilibrium is: CO(g) + H2O(g) ⇌ CO2(g) + H2(g)
a) inject some CO2(g)
b) remove some CO2(g)
c) increase the temperature
d) decrease the pressure by increasing the volume
H = 41.0 kJ
4. Interpret the following graphs in terms of the changes which must have been imposed on
the equilibrium:
A. The equilibrium is: PCl5(g) + 92.5 kJ ⇌ PCl3(g) + Cl2(g)
a)
b)
B. The equilibrium is: H2O(g) + Cl2O(g) ⇌ 2 HOCl(g) + 70 kJ
a)
Chem 30 Equilibrium, Acids & Bases Workbook
b)
5
ICE Tables
1. 1.00 mol of hydrogen gas and 1.00 mol of iodine gas are sealed in a 1.00 L reaction vessel
and allowed to react at 450C. At equilibrium, 1.56 mol of hydrogen iodide gas is present.
Calculate Kc for the reaction.
2. In an experiment, 2.00 mol of H2(g) and 2.00 mol of F2(g) are introduced into a 1.00 L flask
at 500C. After equilibrium was reached, the concentration of HF(g) was 0.500 mol/L.
Calculate the Kc for this reaction at 500C.
3. Phosphorus pentachloride gas decomposes into phosphorus trichloride gas and chlorine gas.
If the [PCl5(g)]i = 8.1 x 10-3 mol/L and the [PCl3(g)]i = 0.298 mol/L, calculate the Kc. The
[Cl2(g)]eq = 2.00 x 103 mol/L.
Kc = 1.32 x 10-5
4. A 1.0 L flask was filled with 2.0 mol of SO2(g) and 2.0 mol of NO2(g) and heated to 250C.
If the Kc is 2.75, calculate the equilibrium concentrations of all species at 250C.
SO2(g)
+
NO2(g)
⇌
SO3(g)
+
NO(g)
0.75 mol/L
0.75 mol/L
1.25 mol/L
1.25 mol/L
5. At a certain temperature, Kc = 2.6  104 for the reaction of H2O with Cl2O. Calculate the
equilibrium concentrations of each species if 10.0 g of H2O and 20.0 g of Cl2O are mixed in a
sealed 2.00 L flask.
H2O(g)
0.2606 mol/L
Cl2O(g)
⇌
0.113635 mol/L
+
2 HOCl(g)
0.00288 mol/L
pH and pOH
1. Use KW to calculate the [H3O+(aq)] in a solution which has a [OH(aq)] of 1.0  1011 mol/L.
(. 1.0 x 10-3 mol/L)
2. Use KW to calculate the [OH(aq)] in a solution which has a [H+(aq)] of 1.0  104 mol/L.
(1.0 x 10-10 mol/L)
3. Determine the pH from the following [H3O+(aq)] or [OH(aq)].
a) [H3O+(aq)] = 1.0  1013 mol/L pH = __________ acidic or basic ____________
b) [H3O+(aq)] = 1  102 mol/L
pH = ___________
acidic or basic ____________
c) [H3O+(aq)] = 106 mol/L
pH = ___________
acidic or basic ____________
d) [OH(aq)] = 0.0010 mol/L
pH = ___________
acidic or basic ____________
e) [OH(aq)] = 1.0  106 mol/L
pH = ___________
acidic or basic ____________
3a. 13 base 3b. 2 acid 3c. 6 acid 3d. 11 base 3e. 8 base
4. Find the pH of a lime that has [H3O+(aq)] = 0.0120 mol/L. (1.921 )
5. Determine the pH of a blood sample with a [OH(aq)] = 2.6  107 mol/L. (7.41 )
6. An ammonia solution has a pOH of 2.92. What is the concentration of hydroxide ions in
solution? (0.00120 mol/L)
Chem 30 Equilibrium, Acids & Bases Workbook
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7. Find the [H3O+(aq)] and pH of a solution made by dissolving 10.0 g of KOH in water to
make 4.00 L of solution. (2.24 x 10-13 mol/L, 12.649)
8. A sodium hydroxide solution is prepared by dissolving 2.50 g to make 2.00 L of solution.
Calculate both the hydroxide ion and hydrogen ion concentrations. (OH = 0.03125 mol/L, H
= 3.2 x 10-13 mol/L)
9. A 0.728 g sample of hydrogen chloride gas is dissolved in 200 mL of water. Calculate both
the hydroxide ion and hydrogen ion concentrations. (H= 0.0998mol/L, OH = 1.002 x 10-13
mol/L)
10. Calculate the pH of a solution made by dissolving 7.50 g of strontium hydroxide to make
500 mL of solution. (13.392)
11. Complete the following table:
[OH(aq)]
2.50 x 10-9
[H3O+(aq)]
4.0 x 106 mol/L
3.2 x 10-10
3.13 x 10-5
5.0 x 10-4
2.0  1011 mol/L
pH
pOH
5.398
8.602
acid
4.50
base
9.50
Acid/Base/Neutral
10 mol/L
5.563
1.36
Brønsted-Lowry Acids and Bases
1. Label each reactant in the following equations as Brønsted-Lowry acid or base:
a)
HSO3(aq)
b)
NH3(aq)
c)
HF (aq)
d)
H2SO3(aq)
+
+
⇌
H2O(l)
H2O(l)
⇌
⇌
HSO3(aq)
+
+
HS (l)
Chem 30 Equilibrium, Acids & Bases Workbook
⇌
7
H3O+(aq)
+
SO32(aq)
NH4+(aq)
+
OH(aq)
F(aq)
+
H2SO3 (aq)
HSO3(aq)
+
H2S (aq)
2. For each of the following reactions, label each reactant as Brønsted-Lowry acid or base and
complete the reaction.
Cl(l)
a)
HNO2 (aq)
+
b)
CH3COOH (aq)
+
c)
H2S (aq)
NO2(aq)
d)
HCO3(aq)
e)
CO32(aq)
+
SO42(aq)
+
⇌
⇌
+
+
S2 (l)
⇌
+
H2O(l)
⇌
+
+
+
⇌
3. For questions 1 and 2, label each product as conjugate acid or conjugate base.
Strengths of Acids
1. Calculate the pH of 500 mL of each of the following acids:
a) 1.0  102 mol/L HCl(aq)
b) 6.00 mol/L HNO2(aq)
c) 1.50 mol/L H2SO3(aq)
d) 6.8  102 mol/L HNO3(aq)
e) 6.3  101 mol/L HF(aq)
2. A 0.80 mol/L solution of an unknown acid, HX(aq), has a pH of 3.75.
a) Calculate the % reaction.
b) Calculate the Ka, the acid ionization constant.
3. Calculate the pH of a solution containing 0.25 mol/L of an acid with an acid ionization
constant (Ka) of 3.2  106 mol/L.
Chem 30 Equilibrium, Acids & Bases Workbook
8
Strengths of Bases
1. Calculate the pH and pOH for each of the following:
a) 1.00 mol/L NaOH(aq)
b) 1.00 mol/L Ca(OH)2(aq)
c) 0.650 mol/L Al(OH)3(aq)
d) A solution made by dissolving 5.82 g of barium hydroxide in 2.00 L of water.
2. Calculate the Kb for each of the following bases at 25C:
a) NO2(aq)
b) F(aq)
c) HSO3(aq)
d) HCO3(aq)
e) OOCCOO2(aq)
3. Calculate the pH of a 13.5 mol/L solution of H2PO4(aq) using the following reaction:
H2PO4(aq) + H2O(l) ⇌ H3PO4(aq) + OH(aq)
Strengths of Acids and Bases – pH Calculations
Calculate the pH for each of the following solutions. Show all work.
1. 0.32 mol/L Mg(OH)2(aq)
2. 6.00 mol/L NH3(aq)
3. 2.0  104 mol/L KHSO4(aq) **red with both litmus
4. 3.0 mol/L H2S(aq)
5. 0.750 mol/L KF(aq)
6. 0.0505 mol/L HI(aq)
7. 16 mol/L CH3COOH(aq)
8. 2.00 mol/L NaCN(aq)
9. 5.7 mol/L Na2HPO4(aq) **blue with both litmus
10. 0.750 mol/L Al(OH)3(aq)
Chem 30 Equilibrium, Acids & Bases Workbook
9
Predicting Acid-Base Reactions
For each of the following problems:
1. Write a reaction equation.
2. Label reactants as acid or base and products as conjugate acid or conjugate base.
3. Use appropriate arrow notation to indicate reaction predomination (forward or reverse).
1. Solutions of Na2SO3(aq) and HF(aq) are mixed in a beaker.
2. A solution of NH4NO3(aq) and a solution of NaCH3COO(aq) are mixed.
3. Sodium benzoate is often used as a preservative. Write the equation for NaC6H5COO(s)
dissolving in a solution of NaHSO4(aq).
4. A household ammonia solution is mixed with a solution of nitrous acid.
5. Nitric acid and potassium hydroxide solutions are mixed.
6. Sodium sulphate is dissolved into a solution of sulphurous acid.
7. Ammonium fluoride is dissolved in water.
8. Benzoic acid can be used as a starting material for the synthesis of aspirin. Write an equation
for the reaction of benzoic acid with water.
9. In solution, the poisonous gas H2S(g), which is sometimes found in natural gas, reacts with
carbonate ions.
10. Rain water in an area near a tar sands plant could dissolve SO2(g). Write the equation that
represents the reaction of this acid solution with sodium sulphate.
11. Vinegar, a dilute ethanoic acid solution, is used to neutralized some spilled lye (sodium
hydroxide)
12. Sodium hydrogen carbonate may be used directly or in gripe water to neutralize excess
stomach acid, HCl(aq).
13. Sodium hydrogen carbonate is mixed with a solution of potassium dihydrogen phosphate.
Chem 30 Equilibrium, Acids & Bases Workbook
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Polyprotic Acids and Bases
For each of the following problems, write all reaction steps and the net reaction. Assume all
reactions are quantitative.
1. Potassium hydroxide is continuously added to a carbonic acid solution.
2. Lithium hydroxide is continuously added to a citric acid solution, C3H4OH(COOH)3(aq).
Citric acid can be abbreviated H3Ct(aq).
3. Potassium hydroxide is continuously added to sulphuric acid until no further reaction occurs.
4. A sodium carbonate solution is titrated with hydrobromic acid.
5. A solution of sodium hydrogen phosphate is titrated with hydroiodic acid.
pH Curves
For each of the following curves state:
1. the titrant
2. what type of sample is being used eg) monoprotic acid, polybasic substance etc.
3. the endpoint pH value(s)
4. potential indicator for each endpoint
5. the equivalence point(s)
Titration Curve
1.
Titration Curve
12 10 pH
8
6
4
2
2.
-
12 10 pH
50
100
150
50
volume of titrant added (mL)
Chem 30 Equilibrium, Acids & Bases Workbook
8
6
4
2
100
150
volume of titrant added (mL)
11
Titration Curve
3.
Titration Curve
12 -
pH
10
8
6
4
2
4.
-
pH
50
100
150
10
8
6
4
2
50
100
150
volume of titrant added (mL)
volume of titrant added (mL)
5.
12 -
Titration Curve
pH
12
10
8
6
4
-
2 50
100
150
volume of titrant added (mL)
Buffers
1. One of the most important buffers in our bodies is the carbonic acid – hydrogen carbonate
ion system.
a) Write the equation for the reaction that occurs when a strong acid is added to this buffer.
b) Write the equation for the reaction that occurs when a strong base is added to this buffer.
2. Another important buffer in our bodies is the dihydrogen phosphate – hydrogen phosphate
ion system.
a) Write the equation for the reaction of this buffer with ammonia.
b) Write the equation for the reaction of this buffer with carbonic acid.
Chem 30 Equilibrium, Acids & Bases Workbook
12