Download T919 Oxidation and Reduction Past Paper Questions

Chemistry 12 (HL)
Unit 4 / Topics 9 + 19 + 13.2
Oxidation and Reduction Reactions
Practice Problems
(including Electrolysis of Aqueous Solutions and Transition Metals)
PAPER 1 QUESTIONS
1.
What are the oxidation numbers of the elements in the compound phosphoric acid,
H3PO4?
Hydrogen
Phosphorus
Oxygen
A.
+1
+1
–2
B.
+1
+5
–2
C.
+3
+1
–4
D.
+3
+5
–8
2.
In which change does nitrogen undergo oxidation?
–
A.
NO2  N2O4
B.
NO3  NO2
–
C.
N2O5  NO3
D.
NH3  N2
3.
Which statement is correct about an oxidizing agent in a chemical reaction?
+
A.
It reacts with oxygen.
B.
It reacts with H ions.
C.
It loses electrons.
D.
It undergoes reduction.
4.
5.
2+
–
What happens to vanadium during the reaction VO (aq)  VO3 (aq)?
A.
It undergoes oxidation and its oxidation number changes from +4 to +5.
B.
It undergoes oxidation and its oxidation number changes from +2 to +4.
C.
It undergoes reduction and its oxidation number changes from +2 to –1.
D.
It undergoes reduction and its oxidation number changes from +4 to +2.
What is the reducing agent in this reaction?
–
+
2+
Cu(s) + 2NO3 (aq) + 4H (aq) → Cu (aq) + 2NO2(g) + 2H2O(l)
A.
6.
Cu(s)
B.
–
NO3 (aq)
C.
Cu
2+
(aq)
D.
+
H (aq)
The following information is given about reactions involving the metals X, Y and Z and
solutions of their sulfates.
X(s) + YSO4(aq)  no reaction
Z(s) + YSO4(aq)  Y(s) + ZSO4(aq)
When the metals are listed in decreasing order of reactivity (most reactive first), what is
the correct order?
A.
C.
Z>Y>X
Y>X>Z
B.
D.
X>Y>Z
Y>Z>X
p. 1
Chemistry 12 (HL)
7.
Which equations represent reactions that occur at room temperature?
–
–
I.
2Br (aq) + Cl2(aq)  2Cl (aq) + Br2(aq)
–
–
II.
2Br (aq) + I2(aq)  2I (aq) + Br2(aq)
–
–
III.
2I (aq) + Cl2(aq)  2Cl (aq) + I2(aq)
A.
8.
C.
C.
D. I, II and III
+
–
H and e on the left
+
–
H on the right and e on the left
B.
D.
+
–
H on the left and e on the right
+
–
H and e on the right
2
B.
4
C.
6
D.
8
2+
Ca (aq) can oxidize Ni(s)
3+
Fe (aq) can oxidize Ni(s)
B.
D.
2+
(aq) can reduce Ca (aq)
3+
2+
Fe (aq) can reduce Ca (aq)
Ni
2+
A particular voltaic cell is made from magnesium and iron half-cells. The overall equation
for the reaction occurring in the cell is
2+
2+
Mg(s) + Fe (aq) → Mg (aq) + Fe(s)
Which statement is correct when the cell produces electricity?
A.
B.
C.
D.
12.
C. II and III only
From the given standard electrode potentials which statement is correct?
2+
–
Ө
Ca (aq) + 2e  Ca(s)
E = –2.87 V
2+
–
Ө
Ni (aq) + 2e  Ni(s)
E = –0.23 V
3+
–
2+
Ө
Fe (aq) + e  Fe (aq)
E = +0.77 V
A.
11.
B. I and III only
+
What is the coefficient for H when the equation below is balanced?
–
+
2+
__Pb(s) + __NO3 (aq) + __H (aq) → __Pb (aq) + __NO(g) + __H2O(l)
A.
10.
I and II only
The unbalanced equation for the conversion of sulfur dioxide to sulfuric acid is given
below.
___SO2 + ___H2O → ___ H2SO4
Which other species are used, and on which side of the equation, to balance it?
A.
9.
Unit 4 / Topics 9 + 19 + 13.2
Magnesium atoms lose electrons.
The mass of the iron electrode decreases.
Electrons flow from the iron half-cell to the magnesium half-cell.
Negative ions flow through the salt bridge from the magnesium half-cell to the iron
half-cell.
Consider the following reaction:
4+
2+
–
+
H2SO3(aq) + Sn (aq) + H2O(l) → Sn (aq) + HSO4 (aq) + 3H (aq)
Which statement is correct?
A.
B.
C.
D.
H2SO3 is the reducing agent because it undergoes reduction.
H2SO3 is the reducing agent because it undergoes oxidation.
4+
Sn is the oxidizing agent because it undergoes oxidation.
4+
Sn is the reducing agent because it undergoes oxidation.
p. 2
Chemistry 12 (HL)
13.
Unit 4 / Topics 9 + 19 + 13.2
Consider the standard electrode potentials of the following reactions.
3+
–
Cr (aq) + 3e →Cr(s)
–0.75 V
2+
–
Cd (aq) + 2e → Cd(s)
–0.40 V
What is the value of the cell potential (in V) for the following reaction?
2+
3+
2Cr(s) + 3Cd (aq) → 2Cr (aq) + 3Cd(s)
A.
–0.35
B.
–1.15
C.
+0.30
D.
+0.35
14.
What happens at the positive electrode in a voltaic cell and in an electrolytic cell?
Voltaic cell
Electrolytic cell
A.
Oxidation
Reduction
B.
Reduction
Oxidation
C.
Oxidation
Oxidation
D.
Reduction
Reduction
15.
Which changes lead to the production of more moles of metal during the electrolysis of a
molten salt?
I.
using a metal ion with a higher charge
II.
increasing the current
III.
using a longer time
A.
16.
B. I and III only
C. II and III only
D. I, II and III
Aqueous solutions of AgNO3, Cu(NO3)2 and Cr(NO3)3 are electrolyzed using the same
quantity of electricity. How do the number of moles of metal formed compare?
A.
C.
17.
I and II only
Ag = Cu = Cr
Ag < Cu < Cr
B.
D.
Ag > Cu > Cr
Cu > Ag > Cr
–
The cyanide ion, CN , can form two complex ions with iron ions. The formulas of these
4–
3–
ions are [Fe(CN)6] and [Fe(CN)6] . What is the oxidation number of iron in the two
complex ions?
4–
3–
[Fe(CN)6]
[Fe(CN)6]
A.
–4
–3
B.
+2
+3
C.
+3
+2
D.
–3
–4
18.
Aqueous solutions containing different concentrations of NaCl were electrolysed using
platinum electrodes. What is the major product at the positive electrode in each case?
–3
–3
0.001 mol dm NaCl(aq)
1.0 mol dm NaCl(aq)
A.
H2
Na
B.
H2
H2
C.
O2
Cl2
D.
Cl2
O2
19.
Which statement is correct about the electrolysis of copper(II) sulfate solution using
graphite electrodes?
A.
A colourless gas is produced at the negative electrode.
B.
The electrolyte does not change colour.
p. 3
Chemistry 12 (HL)
C.
D.
20.
21.
22.
23.
Unit 4 / Topics 9 + 19 + 13.2
The negative electrode decreases in mass.
A colourless gas is produced at the positive electrode.
3
In the electrolysis of acidified water, if 8.4 cm of hydrogen gas is evolved, what volume
of oxygen gas is evolved?
3
3
3
3
A. 4.2 cm
B. 8.4 cm
C. 12.6 cm
D. 16.8 cm
Which salts form coloured solutions when dissolved in water?
I.
FeCl3
II.
NiCl2
III.
ZnCl2
A.
I and II only
B. I and III only
C. II and III only
D. I, II and III
Which combination is correct for the complex ion in [Co(NH3)4(H2O)Cl]Br?
Oxidation state of cobalt
Shape of the complex ion
Overall charge of the
complex ion
A.
+2
Octahedral
+2
B.
+3
Octahedral
–1
C.
+2
Octahedral
+1
D.
+2
Tetrahedral
+1
Which electrons are lost by an atom of iron when it forms the Fe
A.
One s orbital electron and two d orbital electrons
B.
Two s orbital electrons and one d orbital electron
C.
Three s orbital electrons
D.
Three d orbital electrons
3+
ion?
PAPER TWO QUESTIONS
24.
Consider the following redox equation.
2+
–
+
3+
2+
5Fe (aq) +MnO4 (aq) +8H (aq) → 5Fe (aq) + Mn (aq) + 4H2O(l)
(i)
Determine the oxidation numbers for Fe and Mn in the reactants and in the
products.
(2)
(ii)
Based on your answer to (i), deduce which substance is oxidized.
(iii)
The compounds CH3OH and CH2O contain carbon atoms with different oxidation
numbers. Deduce the oxidation numbers and state the kind of chemical change
needed to make CH2O from CH3OH.
(1)
(3)
25.
A part of the reactivity series of metals, in order of decreasing reactivity, is shown below.
magnesium
zinc
iron
lead
copper
p. 4
Chemistry 12 (HL)
Unit 4 / Topics 9 + 19 + 13.2
silver
If a piece of copper metal were placed in separate solutions of silver nitrate and zinc
nitrate
(i)
determine which solution would undergo reaction.
(1)
(ii)
identify the type of chemical change taking place in the copper and write the
half-equation for this change.
(iii)
state, giving a reason, what visible change would take place in the solutions.
(2)
(2)
26.
3+
Iron in food, in the form of Fe , reacts with ascorbic acid (vitamin C), C6H8O6, to form
dehydroascorbic acid, C6H6O6.
(i)
Write an ionic half-equation to show the conversion of ascorbic acid to
dehydroascorbic acid in aqueous solution.
(1)
(ii)
In the other ionic half-equation Fe
3+
is converted to Fe
2+
. Deduce the
3+
overall equation for the reaction between C6H8O6 and Fe .
(1)
27.
(a)
Some standard electrode potentials are shown in Table 14 of the Data Booklet.
(i)
State three conditions under which the hydrogen electrode is assigned a
potential of zero.
(3)
(ii)
Calculate the cell potential of a cell made by connecting standard copper
and zinc electrodes. State the direction of electron flow in the external circuit
when the cell produces current. Outline the changes occurring at the
electrodes and in the solutions during the process.
(5)
(b)
Using information from Table 14, determine whether or not there is a spontaneous
reaction between copper metal and a solution containing hydrogen ions.
(c)
Using information from Table 14, identify a substance that will oxidize bromide ions
but not chloride ions. Explain your choice, and write an equation for the redox
reaction you have chosen.
(2)
(5)
28.
The following are standard electrode potentials.
Half-equation
Zn
Cr
Fe
Sn
2+
3+
2+
2+
(aq) + 2e
(aq) + 3e
(aq) + 2e
(aq) + 2e
–
–
–
–
2+
–
Cu (aq) + 2e
3+
–
Fe (aq) + e
(a)
ο
E /V
Zn(s)
–0.76
Cr(s)
–0.74
Fe(s)
–0.44
Sn(s)
–0.14
Cu(s)
2+
Fe (aq)
+0.34
+0.77
These values were obtained using a standard hydrogen electrode. Describe the
materials and conditions used in the standard hydrogen electrode. (A suitably
labelled diagram is acceptable.)
(5)
p. 5
Chemistry 12 (HL)
(b)
Unit 4 / Topics 9 + 19 + 13.2
Define the term oxidizing agent in terms of electron transfer and identify the
strongest oxidizing agent in the list above.
(2)
(c)
A cell was set up using zinc in zinc sulfate solution and copper in copper(II) sulfate
solution, both solutions being under standard conditions.
(i)
Calculate the cell potential.
(1)
(ii)
Write an equation for the spontaneous cell reaction.
(2)
(d)
Both zinc and tin are used to coat iron to prevent it from rusting. Once the surface
is scratched, oxygen and water containing dissolved ions come into contact with
the iron and the coating metal.
(i)
State and explain whether zinc or tin would be more effective in preventing
iron from rusting under these conditions.
(2)
(ii)
Electroplating may be used to coat one metal with another metal. Identify the
three factors affecting the amount of metal discharged during electroplating.
(3)
(iii)
Explain why electrolysis of aqueous zinc sulfate is not used for coating with
zinc metal.
(2)
(e)
Another cell was set up as shown below.
(i)
Identify the part of the cell labelled Y and outline its function.
(2)
(ii)
Write an equation for the initial reactions at each electrode and hence write
an equation for the cell reaction.
(4)
(iii)
Describe the direction of electron flow in the external circuit.
(1)
(iv)
Calculate the cell potential.
(1)
p. 6
Chemistry 12 (HL)
Unit 4 / Topics 9 + 19 + 13.2
29.
(a)
The apparatus shown above may be used to carry out a redox reaction.
(i)
State the function of the salt bridge.
(1)
(ii)
Write a half-equation for the oxidation reaction.
(1)
(iii)
The above reactions are carried out under standard conditions.
State what the standard conditions are for the cell.
(2)
(iv)
Using the Data Booklet, calculate the cell potential for the above cell.
(2)
(v)
State and explain what happens to the concentration of the copper(II) ions
when the cell is producing an electric current.
(2)
(vi)
State two observations that could be made if the zinc rod were placed in a
solution of copper(II) ions.
(2)
(b)
The standard electrode potentials for three electrode systems are given below.
3+
–
2+
ο
Ti (aq) + e  Ti (aq)
E = –0.37 V
3+
–
2+
ο
Fe (aq) + e  Fe (aq)
E = +0.77 V
4+
–
3+
ο
Ce (aq) + e  Ce (aq)
E = +1.45 V
(i)
Using the data above, deduce which species is the best reducing agent,
giving a reason in terms of electrons for your answer.
(2)
(ii)
Write an equation, including state symbols, for the overall reaction with the
greatest cell potential.
(2)
(iii)
ο
State and explain the sign of ∆G for the reaction in (b) (ii).
(2)
(c)
(i)
State the name of a solution that would produce only hydrogen and oxygen
when electrolyzed using platinum electrodes.
(1)
(ii)
Draw a diagram of apparatus that would allow the gases produced in the
reaction in (c) (i) to be collected separately. Annotate your diagram to show
p. 7
Chemistry 12 (HL)
Unit 4 / Topics 9 + 19 + 13.2
the polarity of each electrode and the names and relative volumes of each
gas.
(3)
30.
(a)
In one experiment involving the electrolysis of molten sodium chloride, 0.1 mol of
chlorine was formed. Deduce, giving a reason, the amount of sodium formed at the
same time.
(2)
(b)
In another experiment involving the electrolysis of molten sodium chloride, the time
of the electrolysis was halved and the current increased from 1 amp to 5 amp,
compared to the experiment in (a). Deduce the amount of chlorine formed, showing
your working.
(2)
31.
(c)
(i)
State the name of a solution that would produce only hydrogen and oxygen
when electrolyzed using platinum electrodes.
(1)
(ii)
Draw a diagram of apparatus that would allow the gases produced in the
reaction in (c) (i) to be collected separately. Annotate your diagram to show
the polarity of each electrode and the names and relative volumes of each
gas.
(3)
32.
(a)
When a concentrated aqueous solution of sodium chloride is electrolyzed using
inert electrodes, a different gas is produced at each electrode.
(i)
Write equations for the oxidation and reduction half-reactions.
Oxidation half-reaction:
Reduction half-reaction:
(2)
(ii)
Explain why sodium is not formed during the electrolysis of aqueous NaCl
solution.
(1)
(b)
33.
Deduce the products formed during the electrolysis of an aqueous solution of
sodium fluoride. Write an equation for the reaction at the positive electrode (the
anode) and give your reasoning. (4)
Consider the transition metal complex, K3[Fe(CN)6].
(i)
Define the term ligand, and identify the ligand in this complex.
(1)
(ii)
Write the full electron configuration and draw the orbital box diagram of iron in its
oxidation state in this complex, and hence, determine the number of unpaired
electrons in this state.
(3)
(iii)
Explain why many transition metal d-block complexes are coloured.
(3)
p. 8
Chemistry 12 (HL)
34.
Unit 4 / Topics 9 + 19 + 13.2
Elements with atomic number 21 to 30 are d-block elements.
(a)
Identify which of these elements are not considered to be typical transition
elements.
(b) Complex ions consist of a central metal ion surrounded by ligands. Define the term
ligand.
(c) Complete the table below to show the oxidation state of the transition element.
(3)
ion
Cr2O7
2–
[CuCl4]
2–
[Fe(H2O)6]
3+
oxidation state
(d)
Identify two transition elements used as catalysts in industrial processes, stating
the process in each case.
(2)
(e)
Apart from the formation of complex ions and apart from their use as catalysts,
state two other properties of transition elements.
OTHER QUESTIONS
35.
36.
a)
Show, using half reactions and cell potential values, that the decomposition of molten
magnesium iodide would only occur in an electrolytic cell.
b)
Identify the charge on the electrode where the metal would form.
An electroplating cell is set up to plate nickel onto a Hello Kitty key chain. A solution of
–3
0.50 mol dm Ni(NO3)2 is used as the electrolyte.
a)
State the material used for the cathode and its charge.
b)
Write the equation for the half reaction that occurs at the cathode.
c)
State the direction of flow of electrons.
d)
Write the equation for the oxidation reaction.
e)
Calculate the mass of nickel plated on Hello Kitty if a current of 40.0 mA is run
–3
through a 0.50 mol dm Ni(NO3)2 solution for 4.00 min.
4
(F = 9.65 x 10 C/mol; q = I t ; ne = q/F)
 turn
p. 9
Chemistry 12 (HL)
f)
Unit 4 / Topics 9 + 19 + 13.2
Predict and explain how the mass of metal plated would change if a 0.50 mol dm
solution of silver nitrate and a silver electrode were used in place of the nickel
electrode and nickel nitrate.
–3
p. 10