Topic 9 Review
1. Consider the overall reaction taking place in a voltaic cell.
Ag2O(s) + Zn(s) + H2O(l) → 2Ag(s) + Zn(OH)2(s)
What is the role of zinc in the cell?
A.
The positive electrode and the oxidizing agent.
B.
The positive electrode and the reducing agent.
C.
The negative electrode and the oxidizing agent.
D.
The negative electrode and the reducing agent.
2. Nitrogen monoxide may be removed from industrial emissions via a reaction with ammonia as shown by
the equation below.
4NH3(g) + 6NO(g) → 5N2(g) + 6H2O(l)
(i)
Deduce the oxidation number of the nitrogen in the reactants and product.
(ii) Deduce the oxidation and reduction half-equations and identify the oxidizing agent for the
reaction.
3. What happens to iodine when iodate ions, IO3–, are converted to iodine molecules, I2?
A.
It undergoes reduction and its oxidation number changes from –1 to 0
B.
It undergoes oxidation and its oxidation number changes from –1 to 0
C.
It undergoes reduction and its oxidation number changes from +5 to 0
D.
It undergoes oxidation and its oxidation number changes from +5 to 0
4. Which definition of oxidation is correct?
A.
Loss of electrons and a decrease in oxidation number
B.
Loss of electrons and an increase in oxidation number
C.
Gain of electrons and a decrease in oxidation number
D.
Gain of electrons and an increase in oxidation number
5. What is the IUPAC name of Fe2O3?
A.
Iron oxide
B.
Iron(II) oxide
C.
Iron(III) oxide
D.
Di-iron trioxide
6. Nitric acid reacts with silver in a redox reaction.
__ Ag(s) + __ NO3–(aq) + ____ → __ Ag+(aq) + __ NO(g) + ____
Using oxidation numbers, deduce the complete balanced equation for the reaction showing all the
reactants and products.
7. Consider the following reaction:
H2SO3(aq) + Sn4+(aq) + H2O(l) → Sn2+(aq) + HSO4–(aq) + 3H+(aq)
Which statement is correct?
A.
H2SO3 is the reducing agent because it undergoes reduction.
B.
H2SO3 is the reducing agent because it undergoes oxidation.
C.
Sn4+ is the oxidizing agent because it undergoes oxidation.
D.
Sn4+ is the reducing agent because it undergoes oxide
8. Which list represents the halogens in increasing order of oxidizing strength (weakest oxidizing agent
first)?
A.
Cl2 I2 Br2
C.
I2 Cl2 Br2
B.
I2 Br2 Cl2
D.
Cl2 Br2 I2
9. The overall reaction that takes place in a voltaic cell is shown below.
Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l)
(i)
Determine the oxidation number of lead in Pb, PbO2 and PbSO4.
(ii)
Deduce the oxidation and reduction half-equations taking place at the negative lead electrode
(anode) and the positive lead(IV) oxide electrode (cathode). Deduce the oxidizing and reducing
agents and state the direction of the electron flow between the electrodes.
(iii) In order to determine the position of three metals in a reactivity series, the metals were placed in
different solutions of metal ions. The table below summarizes whether or not a reaction
occurred.
Ag+(aq)
Ag(s)
Cu(s)
Reaction
Pb(s)
Reaction
Cu2+(aq)
Pb2+(aq)
No reaction
No reaction
No reaction
Reaction
State the equations for the three reactions that take place. Use this information to place the
metals Ag, Cu and Pb in a reactivity series, with the strongest reducing agent first, and explain
your reasoning.
10. A particular voltaic cell is made from magnesium and iron half-cells. The overall equation for the
reaction occurring in the cell is
Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)
Which statement is correct when the cell produces electricity?
A.
Magnesium atoms lose electrons.
B.
The mass of the iron electrode decreases.
C.
Electrons flow from the iron half-cell to the magnesium half-cell.
D.
Negative ions flow through the salt bridge from the magnesium half-cell to the iron half-cell.
11. Draw a labelled diagram of a voltaic cell made from an Fe(s) / Fe2+(aq) half-cell connected to a Cu(s) /
Cu2+(aq) half-cell. In your diagram identify the positive electrode (cathode), the negative electrode
(anode) and the direction of electron flow in the external circuit.
12. What are the features of a standard hydrogen electrode?
I.
A temperature of 298 K
II.
A carbon electrode
III. Hydrogen gas at 1.01 × 105 Pa (1 atm) pressure
A.
I and II only
C.
B.
I and III only
D.
II and III only
I, II and III
13. What conditions are necessary to directly measure a standard electrode potential (EO)?
I.
A half-cell with an electrode in a 1.0 mol dm–3 solution of its ions.
II.
Connection to a standard hydrogen electrode.
III. A voltmeter between half-cells to measure potential difference.
A.
I and II only
C.
II and III only
B.
I and III only
D.
I, II and III
14. An electrochemical cell is made from an iron half-cell connected to a cobalt half-cell:
The standard electrode potential for Fe2+(aq) + 2e–
Fe (s) is –0.45 V. The total cell potential
obtained when the cell is operating under standard conditions is 0.17 V.
Cobalt is produced during the spontaneous reaction.
(i)
Define the term standard electrode potential and state the meaning of the minus sign in the
value of – 0.45 V.
(ii)
Calculate the value for the standard electrode potential for the cobalt half-cell. Calculate the
free-energy for the cell.
(iii)
Determine the mass of cobalt produced if the iron electrode lost 0.056g.
(iv)
Deduce which species acts as the oxidizing agent when the cell is operating.
(v)
Deduce the equation for the spontaneous reaction taking place when the iron half-cell is
connected instead to an aluminium half-cell.
(vi)
Explain the function of the salt bridge in an electrochemical cell.
15. The conditions used in an electrolytic cell can determine the products formed. Draw an electrolytic cell
illustrating the electrolysis of molten nickel(II) bromide, NiBr2. Include in the diagram the direction of
the electron flow, the polarity of electrodes and state the half-equations for the product formed at each
electrode.
16. Solid sodium chloride does not conduct electricity but molten sodium chloride does. Explain this
difference.
Outline what happens in an electrolytic cell during the electrolysis of molten sodium chloride using
inert electrodes. Deduce equations for the reactions occurring at each electrode.
17. Deduce the equations for the formation of the major product at the positive electrode (anode) when the
following aqueous solutions are electrolysed.
• dilute sodium chloride
• concentrated sodium chloride
18. Electrolysis is used in the electroplating of metals. The same amount of current is passed through
separate aqueous solutions of NiSO4, Sn(SO4)2 and Cr2(SO4)3 in separate electrolytic cells for the same
amount of time. State and explain which cell would deposit the greatest amount (in mol) of metal.
Identify the electrode at which the metal is deposited.
How many grams of that metal would be deposited if the cell ran for 1.5 hrs at 3.5 current?
For the Sn(SO4)2 cell, suggest two factors, other than time and current, that would affect the amount of
metal deposited during electroplating.