chem 12 prov exam review - digital...

Chemistry 12

Provincial Exam Review

Prescribed Learning Outcomes

with Selected Questions

from Past Provincial Exams

— Chemistry 12 Provincial Exam Review — Page 1

CHEMISTRY 12 PROVINCIAL EXAM REVIEW

UNIT I REACTION KINETICS

A: REACTION KINETICS (Introduction)

A1. give examples of reactions proceeding at different rates

A2. describe rate in terms of some quantity (produced or consumed) per unit of time

A3. experimentally determine rate of a reaction

A4. identify properties that could be monitored in order to determine a reaction rate

A5. recognize some of the factors that control reaction rates

A6. compare and contrast factors affecting the rates of both homogeneous and heterogeneous reactions

A7. discuss situations in which the rate of reaction must be controlled

1. A1. DIFF RXN RATES 011

Which of the following reactions would have the greatest reaction rate at room temperature?

A. C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g)

B. Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)

C. AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

D. Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + H2O(l) + CO2(g)

2. A2. DESCRIBE RATE 015

Which of the following could be used to describe the rate of a reaction

A. change in time

change in concentration

B. change in mass

change in concentration

C. change in concentration

change in time

D. change in concentration

change in mass

3. A3. DETERMINE RATE 032

Consider the following reaction:

Fe2O3(s) + 2Al(s) → Al2O3(s) + 2Fe(s)

If 0.50 mol of Fe is produced in 10.0 sec, what is the rate of consumption of Fe2O3 in mol/s ?

A. 5.0 x 10–2

mol/s

B. 2.5 x 10–2

mol/s

C. 1.0 x 10–1

mol/s

D. 5.0 mol/s

Chemistry 12 Unit I Reaction Kinetics


4. A3. DETERMINE RATE 033

Consider the following reaction;

3Cu(s) + 8HNO3(aq) → 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)

A piece of copper is added to a nitric acid solution in an open beaker, allowing the NO(g) to escape. The

following data was obtained:

A. Calculate the reaction rate for the time period 2.0 to 6.0 min. (2 marks)

B. Calculate the mass of copper consumed in the first 5 minutes. (3 marks)

5. A4. MONITOR RATE 018


COCl2(g) → CO(g) + Cl2(g)

Which of the following could be used to determine reaction rate in a closed system?

A. a decrease in gas pressure

B. an increase in gas pressure

C. a decrease in the mass of the system

D. an increase in the mass of the system

6. A6. HOMO/HETERO RXNS012

Which of the following does not affect both homogeneous and heterogeneous reaction rates?

A. addition of a catalyst

B. change in temperature

C. change in surface area

D. change in concentration

7. A7. CONTROL RATE 001

Situations exist in everyday life in which chemical reaction rates must be decreased. Describe one specific

situation and state how the decrease could be attained; explain the principle involved. (2 marks)



B: REACTION KINETICS (Collision Theory)

B1. demonstrate an awareness of the following:

• reactions are the result of collisions between reactant particles

• not all collisions are successful

• sufficient kinetic energy (KE) and favourable geometry are required

• to increase the rate of a reaction one must increase the frequency of successful collisions

• energy changes are involved in reactions as bonds are broken and formed

B2. describe the activated complex in terms of its potential energy (PE), stability and structure

B3. define activation energy

B4. describe the relationship between activation energy and rate of reaction

B5. describe the changes in KE and PE as reactant molecules approach each other

B6. draw and label PE diagrams for both exothermic and endothermic reactions, including ∆H, activation

energy and the energy of the activated complex

B7. relate the sign of ∆H to whether the reaction is exothermic or endothermic

B8. write a chemical equation including the energy term (given a ∆H value)and vice versa

B9. describe the role of the following factors in reaction rate:

• nature of reactants

• concentration

• temperature

• surface area

8. B1. COLLISION THEORY021

Which of the following would result in a successful collision between reactant particles?

A. particles have sufficient KE

B. particles convert all their PE into KE

C. particles are in an excited state and are catalyzed

D. particles have sufficient KE and proper molecular orientation

9. B1. COLLISION THEORY023

Using collision theory, explain why reactions between two solutions occur more rapidly than reactions

between two solids. (2 marks)

10. B2. ACTIVATED COMPLX006

An activated complex is a chemical species that is

A. stable and has low PE.

B. stable and has high PE.

C. unstable and has low PE.

D. unstable and has high PE.

11. B3. ACTIVATION ENRGY005

The minimum amount of energy required to overcome the energy barrier in a chemical reaction is the

A. heat of reaction.

B. activation energy.

C. KE of the reactants.

D. enthalpy of the products.



12. B4. EA & RATE 010

A certain reaction is able to proceed by various mechanisms. Each mechanism has a different Ea and

results in a different overall rate. Which of the following best describes the relationship between the Ea

values and the rates?

A.

B.

C.

D.

13. B4. EA & RATE 011

What is the relationship between the activation energy and the rate of a reaction?

A. When the activation energy is high, the rate of reaction is fast.

B. When the activation energy is low, the rate of reaction is slow.

C. When the activation energy is high, the rate of reaction is slow.

D. There is no relationship between activation energy and rate of reaction.

14. B5. PE/KE CHANGES 008

The following diagram shows reactant molecules approaching one another:

What is happening to the kinetic energy and the potential energy?

Kinetic Energy Potential Energy

A. decreasing decreasing

B. decreasing increasing

C. increasing increasing

D. increasing decreasing



15. B6. PE DIAGRAMS 032

Consider the following PE diagram:

The activation energy for the forward reaction is represented by

A. I

B. II

C. III

D. IV

16. B7. ENTHALPY EXO/END004

Consider the reaction:

P2O3(g) + O2(g) → P2O5(g) + 114 kJ

This reaction may be described as

A. exothermic and ∆H = –114 kJ.

B. endothermic and ∆H = 114 kJ.

C. exothermic and ∆H = 114 kJ.

D. endothermic and ∆H = –114 kJ.

17. B8. EQUATIONS/ENERGY005

Which of the following reactions is endothermic?

A. CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) + 890.3 kJ

B. 2Na2O2(s) + 2H2O(l) – 287.0 kJ → 4NaOH(aq) + O2(g)

C. CaO(s) + H2O(l) → Ca(OH)2(aq) ∆H = – 65.2 kJ

D. CaO(s) + 3C(s) → CaC2(s) + CO(g) ∆H = + 464.8 kJ

18. B9. ROLE OF FACTORS 024

Consider the following KE distribution curve for colliding particles:

A. On the diagram above, sketch a line for the distribution of collisions at a higher temperature.

(2 marks)

B. Shade in the area representing the collisions that could result in forming an activated complex at

the lower temperature. (1 mark)



C: REACTION KINETICS (Reaction Mechanisms and Catalysts)

C1. use examples to demonstrate that most reactions involve more than one step

C2. describe a reaction mechanism as the series of steps (collisions) that result in the overall reaction

C3. define catalyst

C4. compare and contrast the PE diagrams for a catalyzed and uncatalyzed reaction in terms of:

• reaction mechanism

• ∆H

• activation energy

C5. identify reactant, product, reaction intermediate and catalyst from a given reaction mechanism

C6. describe the uses of specific catalysts in a variety of situations

19. C2. DESC RXN MECH 018

Consider the following reaction mechanism:

A. Determine the overall reaction. (2 marks)

B. Identify a reaction intermediate. (1 mark)

20. C3. CATALYSTS 014

A substance that increases the rate of a reaction without appearing in the equation for the

overall reaction is a(n)

A. product.

B. catalyst.

C. reactant.

D. intermediate.



21. C4. PE DIAGRAMS CAT 029

Consider the following potential energy diagram for a reaction:

Which of the following represents the correct activation energies?

Forward Catalyzed Reverse Uncatalyzed

Ea Ea

A. 40 kJ 140 kJ

B. 80 kJ 40 kJ

C. 100 kJ 80 kJ

D. 100 kJ 160 kJ

22. C5. RXN MECHANISMS 037

Consider the following proposed reaction mechanism:

Step 1 Fe3+

+ H2O2 → FeH2O23+

Step 2 FeH2O23+

→ FeOH3+

+ HO

Step 3 HO + H2O2 → H2O + HO2

Step 4 FeOH3+

+ HO2 → Fe3+

+ H2O + O2

A. Write the overall reaction. (2 marks)

B. Define the term catalyst and identify a catalyst in the above mechanism. (2 marks)

Chemistry 12 Unit II Dynamic Equilibrium


UNIT II DYNAMIC EQUILIBRIUM

D: DYNAMIC EQUILIBRIUM (Introduction)

D1. describe the reversible nature of most chemical reactions

D2. identify the reversible pathways of a chemical reaction on the PE diagram

D3. relate the changes in rates of the forward and reverse reactions to the changing concentrations of the

reactants and products as equilibrium is established

D4. describe chemical equilibrium as a closed system at constant temperature:

• whose macroscopic properties are constant

• where the forward and reverse reaction rates are equal

• that can be achieved from either direction

• where the concentrations of reactants and products are constant

D5. describe the dynamic nature of chemical equilibrium

D6. infer that a system not at equilibrium will tend to move toward a position of equilibrium

D7. determine entropy and enthalpy changes from a chemical equation (qualitatively)

D8. state that systems tend toward a position of minimum enthalpy and maximum randomness (entropy)

D9. predict the result when enthalpy and entropy factors:

• both favour the products

• both favour the reactants

• oppose one another

1. D3.CHANGE IN RATE/[]029

Consider the following:

2NH3(g) →← N2(g) + 3H2(g)

Initially, NH3 is added to an empty flask. How do the rates of the forward and reverse reactions change as

the system proceeds to equilibrium?

Forward Rate Reverse Rate

A. increases increases

B. increases decreases

C. decreases increases

D. decreases decreases

2. D4.CHARACTER OF EQ 021

Which of the following applies to a chemical equilibrium?

I. Forward and reverse reaction rates are equal

II. Equilibrium can be achieved from either direction

III. Macroscopic properties are constant

A. I only

B. I and II only

C. II and III only

D. I, II and III



3. D5.DYNAMIC EQ 009

A chemical equilibrium is described as “dynamic” because

A. maximum randomness has been achieved.

B. the pressure and temperature do not change.

C. both reactants and products continue to form.

D. the concentrations of chemical species remain constant.

4. D7.ENTROPY/ENTHAPLY 018


2N2(g) + O2(g) + energy →← 2N2O(g)

What positions do minimum enthalpy and maximum entropy tend toward?

Minimum Maximum

Enthalpy Entropy

A. products products

B. products reactants

C. reactants products

D. reactants reactants

5. D8.DRIVING FORCES EQ002

Chemical systems tend to move toward positions of

A. minimum enthalpy and maximum entropy.

B. maximum enthalpy and minimum entropy.

C. minimum enthalpy and minimum entropy.

D. maximum enthalpy and maximum entropy.

6. D9.SPONT/NONSPONT RX017

In which of the following will the driving forces of minimum enthalpy and maximum entropy oppose one

another?

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

B. 2N2(g) + O2(g) → 2N2O(l) ∆H = +164 kJ

C. 2CO(g) + O2(g) → 2CO2(g) ∆H = –566 kJ

D. 4CO2(g) + 6H2O(g) → 2C2H6(g) + 7O2(g) ∆H = +3122 kJ



E: DYNAMIC EQUILIBRIUM (Le Châtelier’s Principle)

E1. describe the term shift as it applies to equilibria

E2. apply Le Châtelier’s principle to the shifting of equilibrium involving the following:

• temperature change

• concentration change

• volume change of gaseous systems

E3. explain the above shifts using the concepts of reaction kinetics

E4. identify the effect of a catalyst on dynamic equilibrium

E5. apply the concept of equilibrium to a commercial or industrial process

7. E2.LE CHATELIER 007

Consider the following equilibrium:

2NO(g) + Br2(g) + energy →← 2NOBr(g)

The equilibrium will shift to the left as a result of

A. adding a catalyst.

B. removing NOBr.

C. increasing the volume.

D. increasing the temperature.

8. E2.LE CHATELIER 024

Consider the following graph for the reaction:

H2(g) + I2(g) →← 2HI(g)

The temperature is increased at t1 and equilibrium is re–established at t2.

A. On the above graph, sketch the line representing the [HI] between time t1 and t3. (1 mark)

B. Calculate the value of Keq after t2. (2 marks)



9. E3.SHIFTS & KINETICS009


PCl3(g) + 3NH3(g) →← P(NH2)3(g) + 3HCl(g)

The volume of the equilibrium system is increased and a new equilibrium is established.

How have the rates been affected?

Rate (forward) Rate (reverse)

A. increased decreased

B. decreased increased

C. decreased decreased

D. did not change did not change

10. E4.CATALYST & EQ 007

A catalyst affects a reversible reaction by

A. making the value of Keq larger.

B. increasing the yield of products.

C. decreasing the value of ∆H for the reaction.

D. enabling equilibrium to achieved more rapidly.

11. E5.APPLY LE CHAT 009

The graph below shows the amount of ammonia produced at various temperatures and pressures during the

Haber process. The reaction is N2(g) + 3H2(g) →← 2NH3(g).

0

20

40

60

80

100

% Yield of NH 3

0 20 40 60 80 100 120

Pressure (x 10 3 kPa)

600 °C

500 °C

400 °C

300 °C

200 °C

industrial

operating

conditions

A. The production of ammonia is an exothermic process. Use information from the graph to support

this statement. (1 mark)

B. Use Le Châtelier’s principle to explain the relationship between increased pressure and the

percentage yield of ammonia. (2 marks)

C. What theoretical conditions produce the greatest yield of ammonia? (1 mark)

D. Industrial operating conditions are indicated on the graph. Explain why these low yield conditions

are used by industry. (2 marks)



F: DYNAMIC EQUILIBRIUM (The Equilibrium Constant)

F1. gather and interpret data on the concentration of reactants and products of a system at equilibrium

F2. write the expression for the equilibrium constant when given the equation for either a homogeneous or

heterogeneous equilibrium system

F3. relate the equilibrium position to the value of Keq and vice versa

F4. predict the effect (or lack of effect) on the value of Keq of changes in the following factors:

• temperature

• pressure

• concentration

• surface area

• catalyst

F5. calculate the value of Keq given the equilibrium concentration of all species

F6. calculate the value of Keq given the initial concentrations of all species and one equilibrium

concentration

F7. calculate the equilibrium concentrations of all species given the value of Keq and the initial

concentrations

F8. determine whether a system is at equilibrium and if not, in which direction it will shift to reach

equilibrium when given a set of concentrations for reactants and products

12. F1.CONC AT EQUIL'M 013


2SO3(g) →← 2SO2(g) + O2(g)

Initially, some SO3 is placed into a 3.0 L container. At equilibrium there is 0.030 mol SO2

present. What is the [O2] at equilibrium?

A. 0.0050 mol/L

B. 0.010 mol/L

C. 0.015 mol/L

D. 0.030 mol/L

13. F2.KEQ EXPRESSIONS 039

Which reaction has the following equilibrium expression?

K eq =[PCl5]

[PCl3][Cl2]

A. PCl3(g) + Cl2(g) →← PCl5(g)

B. PCl3(g) + Cl2(l) →← PCl5(g)

C. PCl5(g) →← PCl3(g) + Cl2(g)

D. PCl5(g) →← PCl3(g) + Cl2(l)



14. F3.POSITION & KEQ 022

An equal number of moles of I2(g) and Br2(g) are placed into a closed container and allowed to establish

equilibrium:

I2(g) + Br2(g) →← 2IBr(g) Keq = 280

Which one of the following relates [IBr] to [I2] at equilibrium?

A. [I2] = [IBr]

B. [I2] < [IBr]

C. [I2] = 2 [IBr]

D. [I2] = 280 [IBr]

15. F4.LE CHAT & KEQ 031


C(s) + 2H2(g) →← CH4(g) ∆H = –74.8 kJ

Which of the following will cause an increase in the value of Keq ?

A. increasing [H2]

B. decreasing the volume

C. finely powdering the C(s)

D. decreasing the temperature

16. F5.KEQ FROM [EQ] 036


N2O4(g) →← 2NO2(g)

An equilibrium mixture contains 4.0 x 10–2

mol N2O4 and 1.5 x 10–2

mol NO2 in a 1.0 L flask.

What is the value of Keq ?

A. 5.6 x 10–3

B. 3.8 x 10–1

C. 7.5 x 10–1

D. 1.8 x 102

17. F6.KEQ FROM [I]/[EQ]031

Consider the following: (4 marks)

H2(g) + I2(g) →← 2HI(g)

Initially, 0.200 mol H2 and 0.200 mol I2 are added to an empty 2.00 L container. At equilibrium,

the [I2] = 0.0200 mol/L. What is the value of Keq ?

18. F7.[EQ]FROM KEQ/[I] 023


H2(g) + Br2(g) →← 2HBr(g) Keq = 12.0

Initially, 0.080 mol H2 and 0.080 mol Br2 are placed into a 4.00 L container. What is the [HBr] at

equilibrium?

19. F8.TRIAL KEQ 026


N2O4(g) →← 2NO2(g) Keq = 9.5 x 10

–3

Initially, 0.060 mol N2O4 and 0.020 mol NO2 are placed into a 2.00 L container. Use calculations to

determine the direction in which the reaction proceeds in order to reach equilibrium.

Chemistry 12 Unit III Solubility Equilibria


UNIT III SOLUBILITY EQUILIBRIA

G: SOLUBILITY EQUILIBRIA (Concept of Solubility)

G1. classify solutions as ionic or molecular given the formula of the solute

G2. describe the conditions necessary to form a saturated solution

G3. describe solubility as the concentration of a substance in a saturated solution

G4. use appropriate units to represent the solubility of substances in aqueous solutions

G5. measure the solubility of a compound in aqueous solution

G6. describe the equilibrium that exists in a saturated aqueous solution

G7. write a net ionic equation that describes a saturated solution

G8. calculate the concentration of the positive and negative ions given the concentration of a solute in an

aqueous solution

1. G1.IONIC/MOLEC SOL'N011

Which of the following will dissolve in water to form an ionic solution?

A. O2

B. CH4

C. NH4Cl

D. CH3OH

2. G1.IONIC/MOLEC SOL'N009

Which of the following dissolves in water to form a molecular solution?

A. KCl

B. Na2O

C. NH4Br

D. C2H5OH

3. G2.SATURATED SOL'NS 006

A saturated solution is formed by adding 10.0 g PbI2(s) to 10.0 mL of water in a beaker. Describe the

situation which exists in the beaker.

A. [Pb2+

] = [I–]

B. moles PbI2(s) = moles Pb2+

(aq)

C. mass of PbI2(s) = mass of Pb2+

(aq)

D. rate of crystalization = rate of dissociation

4. G3.DESCRIBE SOL 006

The solubility of SrCO3 is 2.4 x 10–5

M . How many moles of dissolved solute are present in 100.0 mL of

saturated SrCO3 solution?

A. 5.6 x 10–10

mol

B. 2.4 x 10–6

mol

C. 2.4 x 10–5

mol

D. 2.4 x 10–4

mol

5. G4.UNITS OF SOL 005

Which of the following could be used to express solubility?

A. mol

B. M/s

C. g/mL

D. mL/min



6. G5.MEASURE SOL 006

When 100.0 mL of a saturated solution of BaF2 is heated and all the water is evaporated, 3.6 x 10–4

mol of

solute remains. The solubility of BaF2 is

A. 1.9 x 10–10

M

B. 1.3 x 10–5

M

C. 3.6 x 10–4

M

D. 3.6 x 10–3

M

7. G6.EQUIL & SATURAT'N017

The equation that describes the solubility equilibrium of Ca3(PO4)2 is

A. Ca3(PO4)2(s) →← Ca3

6+(aq) + 2PO4

3–(aq)

B. Ca3(PO4)2(s) →← 3Ca

2+(aq) + 2PO4

3–(aq)

C. Ca3(PO4)2(s) →← 2Ca

3+(aq) + 3PO4

2–(aq)

D. Ca3(PO4)2(s) →← (Ca

2+)3(aq) + (PO4

3–)2(aq)

8. G7.NET IONIC EQN 002

The equation that describes the solubility equilibrium of Ag2CrO4 is

A. Ag2CrO4(s) →← Ag2

2+(aq) + CrO4

2–(aq)

B. Ag2CrO4(s) →← 2Ag

+(aq) + CrO4

2–(aq)

C. Ag2CrO4(s) →← 2Ag(s) + Cr(s) + 2O2(g)

D. Ag2CrO4(s) →← 2Ag

+(aq) + Cr

6+(aq) + 4O

2–(aq)

9. G8.CALCULATE [ION] 021

What are the ion concentrations in 0.30 M CuCl2?

[Cu2+

] [Cl–]

A. 0.10 M 0.20 M

B. 0.20 M 0.10 M

C. 0.30 M 0.30 M

D. 0.30 M 0.60 M



H: SOLUBILITY EQUILIBRIA (Solubility and Precipitation)

H1. describe a compound as having high or low solubility relative to 0.1 M by using a solubility chart

H2. use a solubility chart to predict if a precipitate will form when two solutions are mixed and identify the

precipitate

H3. write a formula equation, complete ionic equation and net ionic equation that represent a precipitation

reaction

H4. use a solubility chart to predict if ions can be separated from solution through precipitation and outline

the process

H5. predict qualitative changes in the solubility equilibrium upon the addition of a common ion

H6. identify an unknown ion through experimentation involving a qualitative analysis scheme

H7. devise a procedure by which the contaminating ions in hard or polluted water can be removed

10. H1.CMPD SOLUBILITY 023

Which of the following has the lowest solubility?

A. CaS

B. CuS

C. FeS

D. MgS

11. H2.PPT 0.1 M SOL'NS 019

When equal volumes of 0.2 M solutions are mixed, which of the following combinations forms a

precipitate?

A. CaS and Sr(OH)2

B. H2SO4 and MgCl2

C. (NH4)2SO4 and K2CO3

D. H2SO3 and NaCH3COO

12. H3.PPT EQUATIONS 016

When equal volumes of 0.20 M Pb(NO3)2 and 0.20 M KCl are mixed, a precipitate of PbCl2 forms.

A. Write the formula equation for the above reaction. (1 mark)

B. Write the complete ionic equation for the above reaction. (1 mark)

C. Write the net ionic equation for the above reaction. (1 mark)

13. H4.SELECTIVE PPT 013

A solution contains both 0.2 M Mg2+

(aq) and 0.2 M Sr2+

(aq) . These ions can be removed separately through

precipitation by adding equal volumes of 0.2 M solutions of

A. OH– and then S

2–

B. Cl– and then OH

–

C. CO32–

and then SO32–

D. SO42–

and then PO43–



14. H5.COMMON ION EFFECT022


CaSO4(s) →← Ca

2+(aq) + SO4

2–(aq)

Which of the following would shift the above equilibrium to the left?

A. adding CaSO4(s)

B. adding MgSO4(s)

C. removing some Ca2+

(aq)

D. removing some SO42–

(aq)

15. H6.QUAL ANALYSIS 008

Consider the following anions:

I. 10.0 mL of 0.20 M Cl–

II. 10.0 mL of 0.20 M OH–

III. 10.0 mL of 0.20 M SO32–

When 10.0 mL of 0.20 M Pb(NO3)2 are added to each of the above, precipitates form in

A. I and II only.

B. I and III only.

C. II and III only.

D. I, II and III.

16. H7.HARD WATER 006

Which of the following could be added to a sample of hard water to remove

both 0.2 M Ca2+

and 0.2 M Mg2+

?

A. 0.2 M S2–

B. 0.2 M Cl–

C. 0.2 M OH–

D. 0.2 M SO42–



I: SOLUBILITY EQUILIBRIA (Quantitative Aspects)

I1. describe the Ksp expression as a specialized Keq expression

I2. write a Ksp expression for a solubility equilibrium

I3. calculate the Ksp for AB and AB2 type compounds when given the solubility of a compound

I4. calculate the solubility of AB and AB2 type compounds from the Ksp

I5. predict the formation of a precipitate by comparing the trial ion product to the Ksp value using specific

data

I6. calculate the maximum concentration of one ion given the Ksp and the concentration of the other ion

I7. demonstrate and describe a method for determining the concentration of a specific ion

17. I2.KSP EXPRESSION 012

The Ksp expression for a saturated solution of Ag2SO3 is

A. Ksp = [2Ag+][SO3

2–]

B. Ksp = [Ag+]

2[SO3

2–]

C. Ksp = [Ag22+

][SO32–

]

D. Ksp = [2Ag+]2[SO3

2–]

18. I3.CALC KSP FROM SOL029

The solubility of CaF2 is 3.3 x 10–4

M. Determine the Ksp value of CaF2.

A. 3.6 x 10–11

B. 1.4 x 10–10

C. 1.1 x 10–7

D. 3.3 x 10–4

19. I3.CALC KSP FROM SOL028

The solubility of CdCO3 is 2.5 x 10

–6 M . Calculate the Ksp value for CdCO3.

A. 6.3 x 10–12

B. 2.5 x 10–6

C. 5.0 x 10–6

D. 1.6 x 10–3

20. I4.CALC SOL FROM KSP023

Calculate the solubility of CaC2O4.

A. 2.3 x 10–9

M

B. 1.2 x 10–5

M

C. 4.8 x 10–5

M

D. 8.3 x 10–4

M

21. I4.CALC SOL FROM KSP013

The solubility of SrF2 is

A. 4.3 x 10–9

M

B. 6.6 x 10–5

M

C. 1.0 x 10–3

M

D. 1.6 x 10–3

M



22. I5.TRIAL ION PRODUCT018

When a solution containing Ag+

is mixed with a solution containing BrO3– , the trial ion product is

determined to be 2.5 x 10–7

. What would be observed?

A. A precipitate would form since trial ion product < Ksp.

B. A precipitate would form since trial ion product > Ksp.

C. A precipitate would not form since trial ion product < Ksp.

D. A precipitate would not form since trial ion product > Ksp.

23. I6.MAX [ION] W/O PPT020

Determine the maximum [Na2CO3] that can exist in 1.0 L of 0.0010 M Ba(NO3)2 without forming a

precipitate.

A. 2.6 x 10–12

M

B. 2.6 x 10–9

M

C. 2.6 x 10–6

M

D. 5.1 x 10–5

M

24. I7.TITRATIONS 006

Consider the following information and accompanying diagram:

In a titration experiment, AgNO3(aq) was used to determine the [Cl–] in a

water sample and the following data were recorded:

[AgNO3] = 0.125 M

Volume of water sample containing Cl– = 20.00 mL

Initial buret reading of AgNO3 = 5.15 mL

Final buret reading of AgNO3 = 37.15 mL

The equation for this reaction is

Ag+

(aq) + Cl–

(aq) → AgCl(s)

Using the above data, determine the [Cl–] in the water sample. (3 marks)

Chemistry 12 Unit IV Acids and Bases


UNIT IV ACIDS AND BASES

J: ACIDS, BASES and SALTS (Properties and Definitions)

J1. identify acids and bases through experimentation

J2. list general properties of acids and bases

J3. write balanced equations representing the neutralization of acids by bases in solution

J4. define Arrhenius acids and bases

J5. write names and formulae of some common acids and bases and outline some of their common

properties, uses and commercial names

J6. define Brönsted–Lowry acids and bases

J7. identify Brönsted–Lowry acids and bases in an equation

J8. write balanced equations representing the reaction of acids or bases with water

J9. identify an H3O+ ion as a protonated H2O molecule that can be represented in shortened form as H

+(aq)

J10. define conjugate acid–base pair

J11. identify the conjugate of a given acid or base

J12. show that in any Brönsted–Lowry acid–base equation there are two conjugate pairs present

1. J1.IDENTIFY A/B 002

Which of the following tests could be used to distinguish between 1.0 M HCl and 1.0 M NaOH?

I. electrical conductivity

II. reaction with zinc to produce hydrogen gas

III. colour of the indicator phenolphthalein

A. III only

B. I and II only

C. II and III only

D. I, II, and III

2. J2.A/B PROPERTIES 031

Which of the following is a property of sodium hydroxide?

A. feels slippery

B. releases H3O+ in aqueous solution

C. changes litmus paper from blue to red

D. reacts with magnesium to produce hydrogen gas

3. J3.NEUTRALIZAT'N EQN005

Which of the following represents the complete neutralization of H3PO4 by NaOH?

A. H3PO4 + NaOH → NaH2PO4 + H2O

B. H3PO4 + 3NaOH → Na3PO4 + 3H2O

C. H3PO4 + 2NaOH → Na2HPO4 + 2H2O

D. H3PO4 + NaOH → NaH + HPO4 + H2O

4. J4.ARRHENIUS A/B 005

An Arrhenius base is defined as a compound that

A. accepts OH– in solution.

B. releases OH– in solution.

C. accepts protons in solution.

D. donates protons in solution.



5. J5.COMMON A/B 002

Caustic soda, NaOH, is found in

A. fertilizers.

B. beverages.

C. toothpaste.

D. oven cleaners.

6. J6.DEFINE B–L A/B 005

A Brønsted–Lowry acid

A. donates protons to a Brønsted–Lowry acid.

B. donates protons to a Brønsted–Lowry base.

C. accepts protons from a Brønsted–Lowry acid.

D. accepts protons from a Brønsted–Lowry base.

7. J7.IDENTIFY B–L A/B 033

Consider the following Brønsted–Lowry equilibrium:

C6H5NH2(aq) + H2O(l) →← C6H5NH3

+(aq) + OH

–(aq)

The substances acting as acids and bases from left to right are

A. acid, base, acid, base.

B. acid, base, base, acid.

C. base, acid, acid, base.

D. base, acid, base, acid.

8. J8.EQNS W/ H2O 004

Water acts as an acid when it reacts with which of the following?

I. CN–

II. NH3

III. HClO4

IV. CH3COO–

A. I and IV only

B. II and III only

C. I, II and IV only

D. II, III and IV only

9. J9.HYDRONIUM ION 001

A hydronium ion has the formula

A. H2+

B. OH–

C. H2O+

D. H3O+

10. J10.DEFINE CONJ A/B 003

A. Define the term Brønsted–Lowry conjugate acid–base pair. (1 mark)

B. Give an example of a conjugate acid–base pair. (1 mark)



11. J11.CONJUGATE A/B 018

What is the conjugate acid and what is the conjugate base of HPO42–

?

Conjugate Acid Conjugate Base

A. PO43–

H2PO4–

B. H2PO4– PO4

3–

C. H2PO4– H3PO4

D. H3PO4 PO43–

12. J12.B–L THEORY 006


HS– + H3BO3

→← H2BO3

– + H2S

The two species acting as Brønsted–Lowry bases in the above equilibrium are

A. HS– and H2S

B. H2BO3 and H2S

C. HS– and H2BO3

–

D. H2BO3 and H2BO3–



K: ACIDS, BASES and SALTS (Strong and Weak Acids and Bases)

K1. relate electrical conductivity in a solution to the concentration of ions

K2. classify an acid or base in solution as either weak or strong by comparing conductivity

K3. define a strong acid and a strong base

K4. define a weak acid and a weak base

K5. write equations to show what happens when strong and weak acids and bases are dissolved in water

(dissociation, ionization)

K6. compare the relative strengths of acids or bases by using a table of relative acid strengths

K7. identify and explain why the strongest acid in aqueous solutions is H3O+ and the strongest base in

aqueous solutions is OH–

K8. predict whether products or reactants are favoured in an acid–base equilibrium by comparing the

strength of the two acids (or two bases)

K9. compare the relative concentrations of H3O+ (or OH

–) between two acids(or between two bases) using

their relative positions on an acid strength table

K10. define amphiprotic

K11. identify chemical species that are amphiprotic

K12. describe situations in which H2O would act as an acid or base

13. K1.CONDUCTIVITY & []011

The electrical conductivities of 0.10 M solutions of NaCl, HCN, and HNO2 are measured. The order by

conductivity from highest to lowest is

A. NaCl > HNO2 > HCN

B. HCN > HNO2 > NaCl

C. NaCl > HCN > HNO2

D. HNO2 > HCN > NaCl

14. K2.CLASSIFY A/B 015

When comparing equal volumes of 0.10 M HNO3 with 0.10 M HNO2, what would be observed?

A. The pH values would be the same.

B. The electrical conductivities would be different.

C. The effects on blue litmus paper would be different.

D. The volumes of 0.10 M NaOH needed for neutralization would be different.

15. K3.DEFINE STRONG A/B002

Which of the following is a property of 1.0 M HCl but not a property of 1.0 M CH3COOH ?

A. turns litmus red

B. ionizes completely

C. has a pH less than 7.0

D. produces H3O+ in solution

16. K4.DEFINE WEAK A/B 003

Which of the following statements applies to 1.0 M NH3(aq) but not to 1.0 M NaOH(aq)?

A. partially ionizes

B. neutralizes an acid

C. has a pH greater than 7

D. turns bromocresol green from yellow to blue



17. K5.A/B EQUATIONS 008

An equation representing the reaction of a weak acid with water is

A. HCl + H2O → H3O+ + Cl

–

B. NH3 + H2O →← NH4+ + OH

–

C. HCO3– + H2O →← H2CO3 + OH

–

D. HCOOH + H2O →← H3O+ + HCOO

–

18. K6.REL STRENGTH A/B 030

List the bases C2O42–

, NH3, and PO43–

in order from strongest to weakest.

A. PO43–

> NH3 > C2O42–

B. C2O42–

> NH3 > PO43–

C. NH3 > PO43–

> C2O42–

D. PO43–

> C2O42–

> NH3

19. K7.STRONGEST AQ A/B 003

In aqueous solutions, H3O+ is the strongest acid present. This phenomenon is called the levelling effect.

Explain why this occurs. (2 marks)

20. K8.POSITION OF EQM 026

Consider the equilibrium:

C6H5COOH + NO2– →← HNO2 + C6H5COO

–

Identify the stronger acid and predict whether reactants or products are favoured.

Stronger Acid Side Favoured

A. HNO2 reactants

B. HNO2 products

C. C6H5COOH reactants

D. C6H5COOH products

21. K9.RELATIVE [H+] 009

A student records the pH of 0.1 M solutions of two acids:

Acids pH

X 4.0

Y 2.0

Which of the following statements can be concluded from the above data?

A. Acid X is stronger than acid Y.

B. Acid X and Y are both weak.

C. Acid X is diprotic while acid Y is monoprotic.

D. Acid X is 100 times more concentrated than acid Y

22. K10.DEF AMPHIPROTIC 005

The term amphiprotic describes a substance that can act as

A. a proton donor and as a proton acceptor.

B. a proton donor but not as a proton acceptor.

C. a proton acceptor but not as a proton donor.

D. neither a proton donor nor as a proton acceptor.



23. K11.ID AMPHIPROTIC 015

Which of the following chemical species are amphiprotic in aqueous solution?

I. F–

II. NH4+

III. HPO42–

A. I only.

B. II only.

C. III only.

D. II and III only.



L: ACIDS, BASES and SALTS (Kw , pH, pOH)

L1. write equations representing the ionization of water using either H3O+ and OH

– or H

+ and OH

–

L2. write the equilibrium expression for the ion product constant of water, Kw

L3. predict the effect of the addition of an acid or base to the equilibrium system:

2H2O(l) →← H3O

+(aq) + OH

– (aq)

L4. state the relative concentrations of H3O+ and OH

–in acid, base and neutral solutions

L5. state the value of Kw at 25°C

L6. describe the variation of the value of Kw with temperature

L7. calculate the concentration of H3O+ (or OH

–) given the other, using Kw

L8. describe the pH scale with reference to everyday solutions

L9. define pH and pOH

L10. define pKw, give its value at 25°C and its relation to pH and pOH

L11. perform calculations relating pH, pOH, H3O+ and OH

–

L12. calculate H3O+ or OH

– from pH and pOH

24. L1.WATER IONIZ'N EQN006

The ionization of water can be represented by

A. 2H2O(l) → 2H2(g) + O2(g)

B. H2O(l) → 2H+

(aq) + O2–

(aq)

C. H2O(l) → H3O+

(aq) + OH–

(aq)

D. 2H2O(l) → H3O+

(aq) + OH–

(aq)

25. L2.KW EXPRESSION 007

Which of the following describes the relationship between [H3O+] and [OH

–] ?

A. [H3O+][OH

–] = 14.00

B. [H3O+] + [OH

–] = 14.00

C. [H3O+][OH

–] = 1.0 x 10

–14

D. [H3O+] + [OH

–] = 1.0 x 10

–14

26. L3.SELF IONIZATION 013

Consider the following equilibrium at 25 °C :

2H2O(l) →← H3O

+(aq) + OH

–(aq)

What happens to [OH–] and pH as 0.1 M HCl is added?

A. [OH–] decreases and pH increases.

B. [OH–] decreases and pH decreases.

C. [OH–] increases and pH increases.

D. [OH–] increases and pH decreases.

27. L4.RELATIVE CONC 007

A basic solution can be defined as one in which

A. [H3O+] is not present

B. [H3O+] is equal to [OH

–]

C. [H3O+] is less than [OH

–]

D. [H3O+] is greater than [OH

–]



28. L5.VALUE OF KW 003

What is the value of the ionization constant for water at 25 °C ?

A. 7.0

B. 14.0

C. 1.0 x 10–7

D. 1.0 x 10–14

29. L6.KW & TEMP 013


2H2O + energy →← H3O+ + OH

–

Which of the following describes the result of decreasing the temperature?

[H3O+] [OH

–] Kw

A. increases increases increases

B. decreases increases decreases

C. increases decreases no change

D. decreases decreases decreases

30. L10.DEFINE PKW 002

Which of the following statements concerning pKw are true?

I. pKw = – log Kw

II. pKw = pH + pOH

III. pKw = [H3O+][OH

–]

A. I and II only

B. I and III only

C. II and III only

D. I, II, and III

31. L11.PH, POH, H+, OH–055

Calculate the pOH of a 0.050 M HBr solution.

A. 0.30

B. 1.30

C. 12.70

D. 13.70

32. L12.CALC [H+] FRM PH015

Determine the pH of 3.0 M KOH .

A. 0.48

B. 11.00

C. 13.52

D. 14.48



M: ACIDS, BASES and SALTS (Ka and Kb Problem Solving)

M1. write Ka and Kb equilibrium expressions

M2. relate the magnitude of Ka or Kb to the strength of the acid or base

M3. given the Ka, Kb and initial concentration, calculate any of the following:

• H3O+

• OH–

• pH

• pOH

M4. calculate the value of Kb for a base given the value of Ka for its conjugate acid (or vice versa)

M5. calculate the value of Ka or Kb given the pH and initial concentration

33. M1.KA & KB EXPRESS'N013

The relationship

H3BO3[ ]OH−[ ]H 2BO 3

−[ ] is the expression for

A. Ka for H3BO3

B. Kb for H3BO3

C. Ka for H2BO3–

D. Kb for H2BO3–

34. M2.MAGNITUDE OF K 013

Four acids are analyzed and their Ka values are determined. Which of the following values represents the

strongest acid?

A. Ka = 2.2 x 10–13

B. Ka = 6.2 x 10–8

C. Ka = 1.7 x 10–5

D. Ka = 1.2 x 10–2

35. M3.KA/KB CALCULAT'N 040

Calculate the pH of 0.35 M H2CO3. (4 marks)

36. M3.KA/KB CALCULAT'N 036

A 0.0200 M solution of methylamine, CH3NH2, has a pH = 11.40. Calculate the Kb for methylamine. (4

marks)

37. M4.CALC KB 026

Calculate the value of Kb for HPO42–

.

A. 4.5 x 10–2

B. 1.6 x 10–7

C. 2.2 x 10–27

D. 6.2 x10–22

38. M5.CALC KA/KB 012

At a particular temperature a 1.0 M H2S solution has a pH = 3.75. Calculate the value of Ka at this

temperature. (4 marks)



N: ACIDS, BASES and SALTS (Hydrolysis of Salts)

N1. write a dissociation equation for a salt in water

N2. write net ionic equations representing the hydrolysis of salts

N3. predict qualitatively whether a salt solution would be acidic, basic, or neutral

N4. determine whether an amphiprotic ion will act as a base or an acid in solution

39. N1.SALT DISSOCIAT'N 006

The dissociation of NH4NO3 is represented by

A. NH4NO3(s) → NH4+

(aq) + NO3–

(aq)

B. NH4+

(aq) + NO3–

(aq) → NH4NO3(s)

C. NH4+

(aq) + H2O(l) → H3O+

(aq) + NH3(aq)

D. NO3–

(aq) + H2O(l) → HNO3(aq) + OH–

(aq)

40. N2.HYDROLYSIS EQN 017

The equation for the predominant hydrolysis of NH4NO3 can be represented by

A. NH4NO3(s) →← NH4

+(aq) + NO3

–(aq)

B. NH4+

(aq) + H2O(l) →← H3O

+(aq) + NH3(aq)

C. NO3–

(aq) + H2O(l) →← HNO3(aq) + OH

–(aq)

D. NH4NO3(aq) + H2O(l) →← H3O

+(aq) + NH3NO3

–(aq)

41. N3.SALT HYDROLYSIS 042

Which of the following salt solutions will be neutral?

A. 1.0 M NH4Cl

B. 1.0 M LiClO4

C. 1.0 M K2C2O4

D. 1.0 M NaHCO3

42. N4.AMPHIPROTIC IONS 013

A solution made from baking soda (NaHCO3) has an amphiprotic anion which is

A. basic since Ka < Kb

B. basic since Ka > Kb

C. acidic since Ka < Kb

D. acidic since Ka > Kb



O: ACIDS, BASES and SALTS (Indicators)

O1. describe an indicator as a mixture of a weak acid and its conjugate base, each with distinguishing

colours

O2. describe the term transition point of an indicator, including the conditions that exist in the equilibrium

system

O3. describe the shift in equilibrium and resulting colour changes as an acid or a base is added to an

indicator

O4. predict the approximate pH at the transition point using the Ka value of an indicator

O5. predict the approximate Ka value for an indicator given the approximate pH range of the colour change

43. O1.INDICATOR COMP 005

A chemical indicator in solution consists of

A. a weak acid and its conjugate acid.

B. a weak acid and its conjugate base.

C. a strong acid and its conjugate acid.

D. a strong acid and its conjugate base.

44. O2.TRANSITION PT 013

Which of the following applies at the transition point for all indicators, HInd ?

A. [HInd] = [Ind–]

B. [Ind–] = [H3O

+]

C. [H3O+] = [OH

–]

D. [HInd] = [H3O+]

45. O3.INDICATOR SHIFT 036

Consider the following equilibrium for the chemical indicator phenol red, HInd, at a pH = 7.3 (orange).

HInd (yellow) + H2O →← H3O+ + Ind

–(red)

When some NaOH is added, what stress is imposed on the equilibrium and what colour change occurs?

Stress Indicator Colour Change

A. increased [H3O+] turns red

B. decreased [H3O+] turns red

C. increased [H3O+] turns yellow

D. decreased [H3O+] turns yellow

46. O4.PH AT TRANSITION 012

A chemical indicator has a Ka = 2.5 x 10–5

. Determine the pH at the transition point.

A. 2.30

B. 4.60

C. 7.00

D. 9.40

47. O5.INDICATOR KA 015

A chemical indicator has a transition point at a pOH = 8.0. Calculate its Ka value and identify the indicator.

A. Ka = 1 x 10–8

, phenol red

B. Ka = 1 x 10–6

, methyl red

C. Ka = 1 x 10–8

, thymol blue

D. Ka = 1 x 10–6

, chlorophenol red



P: ACIDS, BASES and SALTS (Neutralizations of Acids and Bases)

P1. demonstrate an ability to design and perform a neutralization experiment involving the following:

• primary standards

• standardized solutions

• titration curves

• indicators selected so the end point coincides with the equivalence point

P2. calculate from titration data the concentration of an acid or base

P3. calculate the volume of an acid or base of known molarity needed to neutralize a known volume of a

known molarity base or acid

P4. write formula, complete ionic and net ionic neutralization equations for:

• a strong acid by a strong base

• a weak acid by a strong base

• a strong acid by a weak base

P5. calculate the pH of a solution formed when a strong acid is mixed with a strong base

P6. contrast the equivalence point (stoichiometric point) of a strong acid strong base titration with the

equivalence point of a titration involving a weak acid–strong base or strong acid–weak base

48. P1.NEUTRALIZ'N EXPT 012

In acid–base titrations, the solution of known concentration is called a(n)

A. basic solution.

B. acidic solution.

C. standard solution.

D. indicating solution.

49. P2.TITRATION 025

During a titration, 25.0 mL of H3PO4(aq) is completely neutralized by 42.6 mL of 0.20 M NaOH. Calculate

the concentration of the acid.

A. 0.11 M

B. 0.17 M

C. 0.34 M

D. 1.0 M

50. P3.VOL TO END POINT 024

Calculate the volume of 0.300 M HNO3 needed to completely neutralize 25.0 mL of 0.250 M Sr(OH)2.

A. 10.4 mL

B. 15.0 mL

C. 20.8 mL

D. 41.7 mL

51. P4.NEUTRALIZAT'N EQN013

Write the formula equation and the net ionic equation for the reaction between 0.10 M H2SO4 and 0.100 M

Sr(OH)2. (3 marks)

52. P5.PH OF A/B MIXTURE027

Calculate the pH of a solution prepared by mixing 15.0 mL of 0.50 M HCl with 35.0 mL of 1.0 M NaOH.

53. P6.TITRATION CURVES 045

A strong acid–strong base titration has a pH = 7.0 at the equivalence point. A weak acid–strong base

titration has a pH > 7.0 at the equivalence point.

A. What causes the difference in these pH values? (2 marks)

B. Select one indicator which could be used for both titrations. (1 mark)



Q: ACIDS, BASES and SALTS (Buffer Solutions)

Q1. describe the tendency of buffer solutions to resist changes in pH

Q2. describe the composition of an acidic buffer and a basic buffer

Q3. outline a procedure to prepare a buffer solution

Q4. identify the limitations in buffering action

Q5. describe qualitatively how the buffer equilibrium shifts as small quantities of acid or base are added to

the buffer

Q6. describe common buffer systems present in industrial, environmental, or biological systems

54. Q1.PURPOSE OF BUFFER003

All buffer solutions are able to

A. maintain pH at 7.00.

B. neutralize acidic solutions only.

C. maintain a relatively constant pH.

D. keep the pH of a solution at a constant value.

55. Q2.BUFFER COMPOSIT'N023

Equal moles of which of the following chemicals could be used to make a basic buffer solution?

A. HF and NaOH

B. HCl and NaCl

C. KBr and NaNO3

D. NH3 and NH4Cl

56. Q3.PREPARE BUFFER 006

A buffer solution can be prepared by dissolving equal moles of

A. a weak base and a strong base.

B. a weak acid and its conjugate base.

C. a strong base and its conjugate acid.

D. a strong acid and its conjugate base.

57. Q4.BUFFER LIMITAT'N 001

When 10 mL of 0.10 M Sr(OH)2 is added to 20 mL of a solution of 0.10 M CH3COOH and 0.10 M

NaCH3COO, the pH increases greatly. This result occurs because

A. the solution is a buffer.

B. Sr(OH)2 is a strong base.

C. Sr(OH)2 contains a common ion.

D. the amount of OH– exceeds the buffer’s capacity.

58. Q5.BUFFER SHIFT QUAL011

A. Write the net ionic equation that represents the equilibrium that exists in the buffer system

produced when equal volumes of 1.0 M NH3 and 1.0 M NH4Cl are mixed. (1 mark)

B. Explain why the pH of this buffer system changes very little when a small amount of strong base

is added. (2 marks)

59. Q6.COMMON BUFFERS 001

Which of the following pairs of substances form a buffer system for human blood?

A. HCl and Cl–

B. NH3 and NH2–

C. H2CO3 and HCO3–

D. H3C6H5O7 and HC6H5O72–



R: ACIDS, BASES and SALTS (Acid Rain)

R1. write equations representing the formation of acidic solutions or basic solutions from non–metal and

metal oxides

R2. describe the pH conditions required for rain to be called acid rain

R3. relate the pH of normal rain water to the presence of dissolved CO2

R4. describe sources of NOx and SOx

R5. discuss general environmental problems associated with acid rain

60. R1.OXIDES 028

Which of the following equations describes the reaction that occurs when MgO is added to water?

A. MgO + H2O → Mg(OH)2

B. MgO + H2O → MgO2 + H2

C. MgO + H2O → MgH2 + O2

D. 2MgO + 2H2O → 2MgOH + H2 + O2

61. R2.PH OF ACID RAIN 003

The pH of acid rain could be

A. 5.0

B. 7.0

C. 9.0

D. 11.0

62. R3.PH OF RAIN WATER 011

The pH of normal rainwater is

A. less than 7.0 due to dissolved SO2(g)

B. less than 7.0 due to dissolved CO2(g)

C. greater than 7.0 due to dissolved CO2(g)

D. equal to 7.0 due to dissolved N2 and O2

63. R4.ACID RAIN 003

A common source of NO2 is

A. a fuel cell.

B. a lead smelter.

C. an aluminum smelter.

D. an automobile engine.

64. R5.ACID RAIN 001

SO2 is a waste product in some industrial processes. State the environmental problem associated with

SO2(g), write the equation that accounts for this problem, and give one effect on the natural environment. (2

marks)

Chemistry 12 Unit V Electrochemistry


UNIT V ELECTROCHEMISTRY

S: OXIDATION REDUCTION (Introduction)

S1. define and apply the following:

• oxidation–reduction

• oxidizing agent

• reducing agent

• half–reaction

• redox reaction

S2. determine the following:

• the oxidation number of an atom in a chemical species

• the change in oxidation number an atom undergoes when it is oxidized or reduced

• whether an atom has been oxidized or reduced by its change in oxidation number

S3. relate change in oxidation number to gain or loss of electrons

S4. from data for a series of simple redox reactions, create a simple table of reduction half–reactions

S5. identify the relative strengths of oxidizing and reducing agents from their positions on a half–reaction

table

S6. use a table of reduction half–reactions to predict whether a spontaneous redox reaction will occur

between any two species

1. S1.DEFINITIONS 056

Which of the following describes a strong oxidizing agent?

A. a substance which loses electrons readily

B. a substance which gains electrons readily

C. a substance which has a large increase in oxidation number

D. a substance which has a small increase in oxidation number


In which reaction is nitrogen reduced?

A. 2NO + O2 → 2NO2

B. 4NH3 + 5O2 → 4NO + 6H2O

C. Cu2+

+ 2NO2 + 2H2O → Cu + 4H+ + 2NO3

–

D. 4Zn + 10H+ + NO3

– → 4Zn

2+ + NH4

+ + 3H2O


Consider the following redox reaction:

2MnO4– + 3ClO3

– + H2O → 3ClO4

– + 2MnO

2+ 2OH

–

The reducing agent is

A. H2O

B. ClO3–

C. MnO2

D. MnO4–

4. S2.OXIDATION NUMBER 072

What is the oxidation number of S in S2O62–

?

A. +3

B. +5

C. +6

D. +7



5. S3.OX NUM & REDOX 013

The oxidation number of zinc in a reaction increases by 2. This indicates that

A. zinc is reduced and loses 2 electrons.

B. zinc is reduced and gains 2 electrons.

C. zinc is oxidized and loses 2 electrons.

D. zinc is oxidized and gains 2 electrons.

6. S4.HALF–RXN SERIES 022

A solution containing Pd2+

reacts spontaneously with Ga to produce Pd and Ga3+

. However, a solution

containing Pd2+

does not react with Pt. The metals, in order of increasing strength as reducing agents, are

A. Pt < Pd < Ga

B. Pt < Ga < Pd

C. Ga < Pt < Pd

D. Ga < Pd < Pt

7. S5.OA/RA STRENGTHS 034

Which of the following is the weakest oxidizing agent?

A. Cl2

B. Al3+

C. Sn2+

D. acidified Cr2O72–

8. S6.SPONTANEOUS RXNS 042

Which of the following could react spontaneously with Ag metal?

A. Cl–

B. Fe2+

C. acidified SO42–

D. acidified NO3–



T: OXIDATION REDUCTION (Balancing Redox Equations)

T1. balance a half–reaction in solution (acid, base, neutral)

T2. balance a net ionic redox reaction in acid and base solution

T3. write the equations for reduction and oxidation half–reactions given a redox reaction

T4. identify reactants and products for several redox reactions performed in a laboratory and balance the

equations

T5. select a suitable reagent to be used in a redox titration in order to determine the concentration of a

species

T6. determine the concentration of a species by performing a redox titration

9. T1.BALANCE HALF–RXN 028

Which of the following is the balanced half–reaction for

N2O → NH3OH+ (acidic)

A. N2O + 4H+ + 3e

– → NH3OH

+

B. N2O + 3H+ + H2O → NH3OH

+ + 2e

–

C. N2O + 6H+ + H2O → 2NH3OH

+ + 4e

–

D. N2O + 6H+ + H2O + 4e

– → 2NH3OH

+

10. T1.BALANCE HALF–RXN 026

Consider the following half–reaction in a basic solution:

Ag2O3 → AgO (basic)

The balanced half–reaction is

A. Ag2O3 + 4H+ + 4e

– → AgO + 2H2O

B. Ag2O3 + 2H+ + 2e

– → 2AgO + H2O

C. Ag2O3 + H2O + 2e– → 2AgO + 2OH

–

D. Ag2O3 + 2H2O + 4e– → AgO + 4OH

–

11. T2.BALANCE REDOX 024

Balance the following redox equation: (4 marks)

ClO3– + S2O3

2– → S4O6

2– + Cl

– (acidic)

12. T2.BALANCE REDOX 019

Balance the following redox reaction in basic solution. (5 marks)

SO32–

+ MnO4– → SO4

2– + MnO2 (basic)

13. T3.HALF–RXNS 013

Consider the following redox reaction which occurs in a lead–acid storage cell:

PbO2(s) + Pb(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l)

The balanced reduction half–reaction is

A. Pb → Pb2+

+ 2e–

B. Pb + SO42–

→ PbSO4 + 2e–

C. 2H2SO4 + 2Pb + 2e– → 2PbSO4 + 2H2O

D. PbO2 + 4H+ + SO4

2– + 2e

– → PbSO4 + 2H2O



14. T4.LAB REDOX RXNS 010

What occurs when a piece of Zn is placed in 1.0 M Cu(NO3)2 ?

A. [Cu2+

] decreases

B. [Zn2+

] decreases

C. [NO3–] increases

D. no change occurs

15. T5.REAGENT FOR TITRA015

Which of the following could be titrated using acidified MnO4– ions?

A. Na+

B. IO3–

C. SO42–

D. H2O2

16. T6.[] REDOX TITRAT'N018

A titration is performed to determine the concentration of Fe2+

in 25.00 mL of an FeSO4 solution. It

requires 22.52 mL of 0.015 M KMnO4 to reach the equivalence point according to the following equation:

MnO4– + 5Fe

2+ + 8H

+ → Mn

2+ + 5Fe

3+ + 4H2O

Calculate the [Fe2+

] . (4 marks)



U: OXIDATION REDUCTION (Electrochemical Cells)

U1. define, construct and label the parts of an electrochemical cell

U2. identify the half–reactions that take place at each electrode

U3. predict the direction of movement of each type of ion in the cell

U4. predict the direction of flow of electrons in an external circuit

U5. predict which electrode will increase in mass and which will decrease in mass as the cell operates

U6. predict the voltage of the cell when equilibrium is reached

U7. assign voltages to the reduction half–reactions of oxidizing agents by comparison of several cells

U8. describe the significance of the E° of an electrochemical cell

U9. predict the voltage (E° ) of an electrochemical cell using the table of standard reduction half–cells

U10. predict the spontaneity of the forward or reverse reaction from the E° of a redox reaction

U11. describe how electrochemical concepts can be used in various practical applications

U1.PARTS OF ECELL 041

Use the following cell to answer questions 17 and 18.

17. U3.ION MIGRATION 012

Which of the following represents the relationship between [NO3–] and the mass of the Cu electrode in the

complete cell as it operates?

A.

B.

C.

D.



18. U9.PREDICT E° 030

The E° for the above cell is

A. –1.10 Volts

B. –0.42 Volts

C. +0.42 Volts

D. +1.10 Volts

19. U1.PARTS OF ECELL 039

Draw and label an electrochemical cell using a copper anode and having an E° value > 1.00 V. (2 marks)

U5.ELECTRODE CHANGE 022

Use the following diagram to answer questions 20 to 22.

20. U5.ELECTRODE CHANGE 023

In the above electrochemical cell, how do the mass of the anode and the [Ag+] change as the cell operates?

Mass of the Anode [Ag+]

A. decreases increases

B. increases increases

C. decreases decreases

D. no change decreases

21. U2.ELECTRODE RXNS 023

What is the overall cell reaction?

A. 2Ag + Sn2+

→ Sn + 2Ag+

B. 2Ag + Sn → Sn2+

+ 2Ag+

C. 2Ag+ + Sn

2+ → Sn + 2Ag

D. 2Ag+ + Sn → Sn

2+ + 2Ag

22. U9.PREDICT E° 038

What is the value of E° for the cell?

A. –0.94 V

B. –0.66 V

C. +0.66 V

D. +0.94 V



23. U4.ELECTRON FLOW 006

The direction of electron flow in an electrochemical cell is from

A. anode to cathode through the external wire.

B. cathode to anode through the external wire.

C. anode to cathode through the external wire and back through the salt bridge.

D. cathode to anode through the external wire and back through the salt bridge.

U3.ION MIGRATION 009

Use the following diagram to answer questions 24 and 25.

24. U3.ION MIGRATION 010

Which of the following statements apply to this electrochemical cell?

I. Electrons flow through the wire toward the copper electrode.

II. The copper electrode increases in mass.

III. Anions move toward the Zn half cell.

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

25. U6.EQUIL'M VOLTAGE 006

At equilibrium, the voltage of the cell above is

A. –1.10 V

B. 0.00 V

C. +0.42 V

D. +1.10 V

26. U7.REDUCTION POTENTI003

Which of the following statements would be correct if the zinc half–cell had been chosen as the standard

instead of the hydrogen half–cell?

A. The reduction potentials of all half–cells would remain unchanged.

B. The reduction potentials of all half–cells would increase by 0.76 V.

C. The reduction potentials of all half–cells would have positive values.

D. The reduction potential of the hydrogen half–cell would decrease by 0.76 V.



27. U8. SIGNIFICANCE E° 004

Which of the following affects the potentials of electrochemical cells?

I. species used as oxidizing agent

II. temperature

III. concentration of reactants

A. I and II only.

B. II and III only.

C. I and III only.

D. I, II and III.

28. U10.SPONTANEITY E° 008

Consider the following equation:

Cd2+

+ 2I– →← Cd + I2 E°cell = –0.94 V

What is E° for the reduction of Cd2+

?

A. –0.40 V

B. –0.14 V

C. +0.14 V

D. +0.40 V

29. U11.APPLICATIONS 022


I. electrolysis of water

II. electroplating of copper

III. rusting of iron

Which of the above involve non–spontaneous redox reactions?

A. I and II only

B. I and III only

C. II and III only

D. I, II and III



V: OXIDATION REDUCTION (Corrosion)

V1. describe the conditions necessary for corrosion to occur

V2. analyse the process of metal corrosion in electrochemical terms

V3. suggest several methods of preventing or inhibiting corrosion of a metal

V4. describe and explain the principle of cathodic protection

30. V1.CORROSION 007

Which of the following must be present to produce rust by the corrosion of iron?

I. water

II. oxygen

III. salt

A. I only

B. II only

C. I and II only

D. I, II and III

31. V2.METAL CORROSION 005

What happens to iron as it corrodes?

A. It loses electrons and is reduced.

B. It gains electrons and is reduced.

C. It loses electrons and is oxidized.

D. It gains electrons and is oxidized.

32. V3.PREVENT CORROSION012

Describe two chemically different methods of preventing the corrosion of iron. Explain how each method

works. (3 marks)

33. V4.CATHODIC PROTECT 014

Which of the following metals could be used to cathodically protect iron?

A. tin

B. lead

C. zinc

D. copper



W: OXIDATION REDUCTION (Electrolytic Cells)

W1. define electrolysis and electrolytic cell

W2. design and label the parts of an electrolytic cell capable of electrolyzing an aqueous salt (use of over–

potential effect not required)

W3. predict the direction of flow of all ions in the cell

W4. write the half–reaction occurring at each electrode

W5. demonstrate the principles involved in simple electroplating

W6. construct an electrolytic cell capable of electroplating an object

W7. describe the electrolytic aspects of metal refining processes

W8. draw and label the parts of an electrolytic cell used for electrolysis of a molten binary salt

34. W1.DEFINE ELY 006

The process of applying an electric current through a cell to produce a chemical change is called

A. corrosion.

B. ionization.

C. hydrolysis.

D. electrolysis.

35. W2.PARTS OF ELY–CELL010

Consider the following operating cell:

Which of the following describes the cell above?

Electrode #1 Electrode #2 Gas Produced

A. anode cathode H2(g)

B. anode cathode O2(g)

C. cathode anode H2(g)

D. cathode anode O2(g)



W4.ELECTRODE RXNS 045

Use the following diagram to answer questions 36 and 37.

36. W4.ELECTRODE RXNS 046

What reaction occurs at the cathode?

A. 2I– → I2 + 2e

–

B. Cu2+

+ 2e– → Cu

C. H2O → 1/2O2 + 2H

+ + 2e

–

D. 2H2O + 2e– → H2 + 2OH

–

37. W3.ION MIGRATION 005

What happens to the [I–] in the operating cell?

A. [I–] increases overall.

B. [I–] decreases overall.

C. [I–] remains constant overall.

D. [I–] decreases near the anode and increases near the cathode.

38. W5.ELECTROPLATING 010

Which of the following are necessary for electroplating to occur using an electrolytic cell?

I. Two electrodes

II. A metal ion being reduced

III. A direct current power source

A. I and II only.

B. I and III only.

C. II and III only.

D. I, II, and III.

39. W6.ELECTROPLATE CELL008

Draw and label a simple electrolytic cell capable of electroplating an inert electrode with silver. (2 marks)

40. W7.METAL REFINING 005

Draw a diagram of an operating electrolytic cell used to extract pure lead from an impure lead sample.

Identify the electrolyte and the material used for the anode. (3 marks)

41. W8.ELY BINARY SALT 007

A. Draw and label the parts of an electrolytic cell that can be used for the electrolysis of molten NaCl

(1.5 marks)

B. Write the half–reaction that occurs at the cathode. (1 mark)

C. Identify the product at the anode. (0.5 mark)

Chemistry 12 Answers to Questions


ANSWERS TO QUESTIONS

UNIT I REACTION KINETICS

Q SOURCE ANS COMMENTS

1. January - June 2003 C

2. August 2002 C

3. January - June 2003 B

4. January - June 2003 � WR A.

195.45 g - 188.15 g

6.0 min - 2.0 min = 1.83 g/min

B. 200.00 – 189.90 g NO

30.0 g/mol x

3 Cu

2 NO x

63.5 g

mol = 32.1 g

5. January - August 2001 B

6. January - June 2003 C

7. WR Food spoilage → refrigerate

Body enzymes → catalysts

8. January - June 2003 D

9. January - June 2003 � WR Solutions react more quickly than solid because the particles can mix

thoroughly (more surface area) and the particles have greater mobility

(greater number of collisions).

10. Jan - Aug 1999 D

11. Jan - Aug 1999 B

12. Jan - Aug 1999 C

13. January - August 2001 C



16. A

17. January - August 2001 D

18. Jan - Aug 1999 WR

19. January - June 2003 � WR A. 2NO + 2H2 → N2 + 2H2O

B. N2O2 or N2O


21. January - June 2003 A

22. January - August 2001 WR A. 2H2O2 → 2H2O + O2

B. A catalyst is a substance that speeds up a chemical reaction by

providing a lower energy pathway.

Catalyst in mechnism is Fe3+



UNIT II DYNAMIC EQUILIBRIUM






5. April 1997 A



8. August 2001 � A.

B. Keq = [HI]2

[H2][I2] =

(1.0)2

(0.5)(0.1) = 20


10. D

11. June 1995 Scholarship WR A. As the temperature increases the % yield of NH3 decreases. This

indicates a shift to the left. When energy is on the product side, a shift

to the left will occur when the temperature is increased. If energy is on

the product side the reaction is exothermic.

B. At 600 °C when pressure is increased, the % yield of NH3 also

increases. According to Le Chatelier’s principle, when the pressure on

an equilibrium is increased, there will be a shift to the side that

produces fewer gas molecules. In this equilibrium, there would be a

shift to produce more NH3 when the pressure is increased.

C. Low temperature and high pressure. (200 °C and 100 kPa)

D. Higher temperature → faster reaction rate even though reduced yield.

Lower pressure → safer and easier to attain.

12. January - August 2002 A


14. June 1994 Provincial B





17. January - August 2002 WR

H2 + I2 →← 2HI

0.100 M 0.100 M 0.00 M

– x – x +2x

0.1 – x 0.1 – x 2x

0.0200 M 0.0200 M 0.160

Keq = [HI]2

[H2][I2] =

(0.160)2

(0.0200)2 = 64.0


H2 + Br2 →← 2HBr

0.020 M 0.020 M 0.00 M

– x – x +2x

0.020 – x 0.020 – x 2x

Keq = [HBr]2

[H2][Br2] ⇒

(2x)2

(0.020–x)2 = 12.0

x = 0.0127 M

[HBr] = 2x = 0.0254 M

19. January - August 2002 WR Q =

[NO2]2

[N2O4] =

(0.010)2

(0.030) = 0.0033

Q (0.0033) < Keq (0.0095) Equilibrium shifts right, forward direction.



UNIT III SOLUBILITY EQUILIBRIA



2. April - June 2001 D

3. April - June 2001 D





8. April - June 2001 B




12. January - August 2002 WR A. Pb(NO3)2(aq) + 2KCl(aq) → 2KNO3(aq) + PbCl2(s)

B. Pb2+ + 2NO3– + 2K+ + 2Cl– → 2K+ + 2NO3

– + PbCl2(s)

C. Pb2+ + 2Cl– → PbCl2(s)



15. Aug 1998 – Aug 1999 D






21. Aug 1998 - Aug 1999 C



24. January - August 2002 WR mol AgNO3 = 0.03200 L x 0.125 M = 0.00400 mol Ag+

mol of Cl– = 0.00400 mol Ag+ x 1 Cl–

1 Ag+ = 0.00400 mol Cl–

[Cl–] = 0.00400 mol

0.02000 L = 0.200 M



UNIT IV ACIDS AND BASES


1. January 2001 C



4. January 2001 B

5. January 1995 Provincial D

6. April 1994 B



9. April 2001 - August 2001 D

10. January - August 2000 WR A. Bronsted–Lowry conjugate acid–base pairs are two substances that

differ from each other by one proton (H+).

B. H3PO4 and H4PO4–

11. April 2001 - August 2001 B


13. April 2001 - August 2001 A


15. April 2001 - August 2001 B

16. January 2001 A



19. January - August 1999 WR H3O+ is the strongest acid that can exist in water because any acid

stronger will completely dissociate to form H3O+ ions.


21. January 1995 Provincial B

22. A

23. April 2001 – Aug 2001 C







30. August 1996 Provincial A








H2CO3 + H2O →← HCO3– + H3O

+

0.35 M — 0.00 M 0.00 M

– x — +x +x

0.35 – x

≈ 0.35

— x x

Ka = [HCO3

–][H3O+]

[H2CO3] =

x2

0.35 = 4.3 x 10–7

x = 3.9 x 10–4 M

pH = –log (3.9 x 10–4) = 3.41


CH3NH2 + H2O →← CH3NH3+ + OH–

0.0200 M — 0.00 M 0.00 M

– x — +x +x

0.0200 – x — x x

0.0175 — 0.00251 M 0.00251 M

pOH = 14.00 – 11.40 = 2.60

[OH–] = antilog (–2.60) = 0.00251 M

Kb = [CH3NH3

+][ OH–]

[ CH3NH2] =

(0.00251)2

0.0175 = 3.6 x 10–4



H2S + H2O →← HS– + H3O+

1.0 M — 0.00 M 0.00 M

– x — +x +x

1.0 – x — x x

≈ 1.0 — 1.79 x 10–4 1.79 x 10–4

[H3O+] = antilog (–3.75) = 1.79 x 10–4 M

Ka = [HS–][ H3O

+]

[ H2S] =

(1.79 x 10–4)2

1.0 = 3.2 x 10–8















51. January - August 2000 WR 2H+ + SO42– + Sr2+ + 2OH– → 2H2O(l) + SrSO4(s)

52. April - August 2001 � WR [H3O

+] = (15.0 mL)(0.50 M)

(50.0 mL) = 0.15 M

[OH–] = (35.0 mL)(1.0 M)

(50.0 mL) = 0.70 M

[OH–] = 0.70 – 0.15 = 0.55 M

pOH = – log (0.55) = 0.26

pH = 14.00 – 0.26 = 13.74

53. January - August 2002 WR A. When a strong acid is titrated with a strong base, the salt which is

formed in a neutral salt so the pH of the solution is 7.00. When a weak

acid is titrated with a strong base, the salt which is formed in a basic

salt so the pH of the solution is > 7.00.

B. Phenolphthalein

54. January 1993 C



57. D

58. WR A. NH3 + H2O

∅♦ NH4

+ + OH–

B. When a strong base is added to the buffer, the conjugate acid, NH4+,

reacts with the excess OH– and the equilibrium shifts left to compensate

and returns the [OH–] to a level close to what it was initially.

59. August 1996 Provincial C

60. April 2001 - Aug 2001 A

61. April 2001 - Aug 2001 A



64. WR SO2 dissolves into the rain clouds to produce acid rain which then has

environmental impacts such as destroying plant and animal habitats by

upsetting the pH balance.

SO2 + H2O → H2SO3



UNIT V ELECTROCHEMISTRY


1. January - April 2002 B



4. June - August 2002 B


6. January - April 2002 A


8. June - August 2002 D

9. January - April 2002 D


11. June - August 2002 WR 5+ 2+ 2.5+ 1–

6H+ + ClO3– + 6S2O3

2– → 3S4O62– + Cl– + 3H2O

12. January - August 2001 WR 4+ 7+ 6+ 4+

2H+ + 3SO32– + 2MnO4

– → 3SO42– + 2MnO2 + H2O




16. January - April 2002 WR 0.02252 L x 0.015 M = 3.38 x 10–4 mol

3.38 x 10–4 mol MnO4– x

5Fe2+

MnO4– = 1.69 x 10–3 mol Fe2+

[Fe2+] = 1.69 x 10–3 mol

0.0250 L = 0.068 M



19. January - August 2000 WR Cu → Cu2+ + 2e– –0.34

Au3+ + 3e– → Au +1.50

20. June - August 2002 C


Cu

Au

1 M NaNO3

1 M Au(NO3)3(aq) 1 M Cu(NO3)2(aq)




23. August 1996 A



26. April 1996 B


28. June - August 2002 A

29. January - April 2002 A


31. June - August 2002 C

32. June - August 2002 WR 1. Paint the iron → prevents H2O and O2 from coming in contact with Fe.

2. Cathodic protection → attaching a stronger reducing agent such as Zn.

Stronger reducing agent will oxidize first leaving Fe intact.

33. January - April 2002 C







40. June - August 2002 WR

D.C

. + –

1 M AgNO3(aq)

Ag(s) Inert Electrode

D.C

. + –

1 M Pb(NO3)2(aq)

Impure Lead Pure Lead



41. WR A.

B. Na+ + e– → Na(s)

C. 2Cl– → Cl2(g) + 2e– (anode reaction)

Cl2(g) is produced at the anode.

D.C

. + –

NaCl(l)

Inert Electrode

(Pt) Inert Electrode

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