Oxidation and Reduction Reactions

Reaction Types Review

• Decomposition reactions: $CuCO<em>3 + heat \rightarrow CuO + CO</em>2$
• Precipitation reactions: $AgNO<em>3 + NaCl \rightarrow AgCl + NaNO</em>3$ which simplifies to $Ag^+ + Cl^- \rightarrow AgCl(s)$
• Acid-base reactions: $HCl + NaOH \rightarrow H<em>2O + NaCl$ which simplifies to $H^+ + OH^- \rightarrow H</em>2O$

Oxidation-Reduction (Redox) Reactions

• Involve electron transfer.
• Must involve simultaneous oxidation and reduction.
• Require a species to lose electrons (be oxidised).
• Require a species to gain electrons (be reduced).
• Remember: In order for a substance to gain electrons (be reduced) another substance must lose electrons (be oxidised).

Examples

• Burning magnesium: $2Mg + O_2 \rightarrow 2MgO$ (white).
  • $2Mg \rightarrow 2Mg^{2+} + 4e^-$ (Magnesium loses electrons - oxidation).
  • $O_2 + 4e^- \rightarrow 2O^{2-}$ (Oxygen gains electrons - reduction).
  • Magnesium (Mg) has an oxidation state change from (0) to (+2), and Oxygen (O) changes from (0) to (-2).

*Another Example:

$H<em>2 + O</em>2 \rightarrow H_2O$

• Hydrogen is oxidized, losing electrons
• Oxygen is reduced, gaining electrons.

Demonstration

• Zinc reacting with copper sulfate solution: $Zn + CuSO_4$

Reaction

• $Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$
  • Zinc (Zn) is oxidised, losing 2 electrons.
  • Copper ($Cu^{2+}$) is reduced.
  • Oxidation states: Zn (0) -> (+2), Cu (+2) -> (0).

Oxidising and Reducing Agents

• Reducing agent: A species that is oxidised and brings about the reduction of another species.
• Oxidising agent: A species that is reduced and brings about the oxidation of another species.

Example

• $Zn + I<em>2 \rightarrow ZnI</em>2$
  • Zinc loses electrons, is oxidised, and acts as a reducing agent.
  • Iodine accepts electrons, is reduced, and acts as an oxidising agent.
  • Oxidation states: Zn (0) -> (+2), I (0) -> (-1).

Identifying Oxidised and Reduced Species

• Examples:

  • (a) $Zn + 2H^+ \rightarrow Zn^{2+} + H_2$

    • Zn is oxidised, $H^+$ is reduced.

  • (b) $Mg + S \rightarrow MgS$

    • Mg is oxidised, S is reduced.

  • (c) $Cu^{2+} + Fe \rightarrow Cu + Fe^{2+}$

    • $Cu^{2+}$ is reduced, Fe is oxidised.

  • (d) $Cl<em>2 + 2Br^- \rightarrow Br</em>2 + 2Cl^-$

    • $Cl_2$ is reduced, $Br^-$ is oxidised.

Oxidation Number (O.N.)

• The oxidation number of an atom is an arbitrary charge assigned to the atom according to a set of rules.

Rules for Assigning Oxidation Numbers

1. Elements in elemental state have O.N. = 0.

2. For monatomic ions O.N. = charge on ion.

3. For combined oxygen O.N. = -2 (except peroxides, -1).

4. For combined hydrogen O.N. = +1 (except hydrides, -1).

5. For polyatomic species the sum of the O.N. of all the atoms equals the charge on the ion or molecule.

Examples

1. Elements: $O<em>2$, Na, $Cl</em>2$, Fe, Ar, $H_2$ (All neutral elements have O.N. = 0).

2. Monoatomic ions:

   • $Cu^{2+}$ (O.N. = +2)
   • $Cl^-$ (O.N. = -1)
   • $P^{3-}$ (O.N. = -3)

3. Combined oxygen and/or hydrogen:

   • $H<em>2SO</em>4$: 2 (+1) + (S) + 4 (-2) = 0. Therefore, O.N. (S) = +6.

4. Polyatomic ions:

   • $MnO_4^-$: (Mn) + 4 (-2) = -1. Therefore, O.N. (Mn) = +7.

Oxidation States of Chlorine

• $Cl_2$ (chlorine): 0
• HCl (hydrochloric acid): -1
• HClO (hypochlorous acid): +1
• $HClO_2$ (chlorous acid): +3
• $HClO_3$ (chloric acid): +5
• $HClO_4$ (perchloric acid): +7

Using Oxidation Numbers to Identify Redox Reactions

• Oxidation: Increase in O.N.
• Reduction: Decrease in O.N.

Example

• $2Mg + O_2 \rightarrow 2MgO$
  • Mg: (0) -> (+2) (oxidised)
  • O: (0) -> (-2) (reduced)

Copper with Nitric Acid Reaction

• $Cu(s) + 4H^+(aq) + 2NO<em>3^-(aq) \rightarrow Cu^{2+}(aq) + 2NO</em>2(g) + 2H_2O(l)$
  • Oxidation states: Cu (0) -> (+2), N (+5) -> (+4).
  • Copper is oxidised, and nitrogen is reduced.

Volcano Demonstration

• $(NH<em>4)</em>2Cr<em>2O</em>7(s) \rightarrow Cr<em>2O</em>3(s) + N<em>2(g) + 4H</em>2O(g)$

• $Cr<em>2O</em>7^{2-} \rightarrow Cr^{3+}$
  *Determining Oxidation Numbers

• $(NH<em>4)</em>2Cr<em>2O</em>7(s) \rightarrow Cr<em>2O</em>3(s) + N<em>2(g) + 4H</em>2O(g)$

• $(-3)(+1) (+6) (-2) (+3) (-2) (0) (+1) (-2)$

*N in $(NH<em>4)</em>2Cr<em>2O</em>7(s)$ goes from -3 to 0 -> oxidation

*Cr in $(NH<em>4)</em>2Cr<em>2O</em>7(s)$ goes from +6 to +3 -> reduction

*N is the reducing agent and Cr is the oxidising agent

Predicting Reaction Tendency

• Why does zinc strip react in copper sulfate solution, but copper strip does not react in zinc sulfate solution?
• $Zn + CuSO_4$: Reaction occurs.
• $Cu + ZnSO_4$: No reaction.

Standard Reduction Potentials

• Table lists half-equations (written as reductions) in terms of decreasing tendency to be reduced.

  • $F_2$: Greatest tendency to be reduced (best oxidising agent) - top left.
  • Li: Greatest tendency to be oxidised (best reducing agent) - bottom right.

• $E^o$ (volts) is a measure of the tendency for the reduction half-equations to occur.
  *Half reaction $E^o$ electrode potential, V

• $Ag^+ + e^- \rightarrow Ag(s)$ +0.80

• $Cu^{2+} + 2e^- \rightarrow Cu(s)$ +0.34

• $2H^+ + 2e^- \rightarrow H_2(g)$ 0.00

• $Cd^{2+} + 2e^- \rightarrow Cd(s)$ - 0.40

• $Zn^{2+} + 2e^- \rightarrow Zn(s)$ - 0.76

Steps to Predict Reaction Tendency

1. Identify the two half-equations.
2. Add the $E^o$s (after changing the sign for the oxidation reaction) - the sum is the cell voltage.
3. A positive $E^o$ (cell voltage) indicates the reaction is possible (spontaneous).

Example 1

• Predict what would happen when zinc metal is placed in a solution containing $Cu^{2+}$ ions.
• Two reduction half reactions from the table:
  • $Cu^{2+} + 2e^- Cu(s)$ $+0.34 V$
  • $Zn^{2+} + 2e^- Zn(s)$ $- 0.76 V$

*Re-write as oxidation and reduction half reactions:
* $Cu^{2+} + 2e^- Cu(s)$ $+0.34 V$
* $Zn(s) Zn^{2+} + 2e^- +0.76 V$

• Thus, the redox reaction is:
  $Cu^{2+} + Zn(s) Cu(s) + Zn^{2+} + 1.10 V$

  • Zinc dissolves and Cu metal deposits.
  • Zinc is oxidised.
  • Copper is reduced.

Example 2

• Predict what would happen when copper metal is placed in a solution containing $Zn^{2+}$ ions.
  *Two reduction half reactions from the table:

  • $Cu^{2+} + 2e^- Cu(s) +0.34 V$
  • $Zn^{2+} + 2e^- Zn(s) - 0.76 V$
    *Re-write as oxidation and reduction half reactions:
  • $Cu(s) Cu^{2+} + 2e^- -0.34 V$
  • $Zn^{2+} + 2e^- Zn(s) -0.76 V$

• Thus, the redox reaction is:
  $Zn^{2+} + Cu(s) Zn(s) + Cu^{2+} -1.10 V$

  • No reaction, since potential is negative.

Examples of Redox Reactions in Natural Systems

• Burning $CH<em>4 + O</em>2 \rightarrow CO<em>2 + H</em>2O$
• Cell respiration: $C<em>6H</em>{12}O<em>6 + O</em>2 \rightarrow CO<em>2 + H</em>2O$

Redox Reactions in Biological Systems

• Oxidation-reduction reactions are very important in biological systems surrounding the ATP-PC, Glycolysis, Krebs Cycle, and Electron Transport Chain.
• Energy is necessary for muscle contraction, which is critical for human movement.
• Chemical substances such as carbohydrates, lipids, and proteins provide the energy for movement by muscle contraction.
• The nutrients yield usable forms of energy (a process called catabolism) from the chemical activities (i.e. metabolism) that take place in the muscle fibers.
• Since the body’s metabolism is increased with exercise, the oxidation or breakdown of the nutrients is increased to provide more energy in the form of adenosine triphosphate, ATP.
• The catabolism of glucose by way of Glycolysis and the Krebs cycle supplies the electrons that produce energy within the Electron Transport Chain.

Electron Transport Chain

• Includes reactions involving NADH, FADH2, Coenzyme Q, Cytochromes, and oxygen to produce ATP and water.
  *NADH + $H^+$ FADH2 FAD Oxidized CoenzymeQ Reduced NAD+
  *$Fe^{+2}$ Cytochrome b $Fe^{+3}$ 2e- ADP + Pi ATP
  *$Fe^{+2}$ Cytochrome c1 and c $Fe^{+3}$
  *$Fe^{+3}$ Cytochrome a $Fe^{+2}$ 2e- ADP + Pi ATP
  *$Fe^{+2}$ Cytochrome a3 $Fe^{+3}$ 2e- ADP + Pi ATP
• $2e^-$ 2$H^+$ + ½ $O<em>2 H</em>2O$

Chemical Change and Electricity

• Electrochemical cells use a chemical reaction to generate electricity (i.e., an emf).
• Need to separate the oxidation process (anode) and reduction process (cathode) in order to harness electrical energy.
• Examples: Torch battery (dry cell) and car battery (lead-acid accumulator).

Electrochemical Cells

• Anode: Oxidation (e.g., $Zn \rightarrow Zn^{2+} + 2e^-$
• Cathode: Reduction (e.g., $Cu^{2+} + 2e^- \rightarrow Cu$).
• Electron flow from anode to cathode through the circuit.
• Salt bridge maintains charge balance.

Electrolytic Cells

• Use electricity to produce chemical change.
• Examples: Metal extraction, electrorefining, electroplating.
• Anode: Oxidation.
• Cathode: Reduction.
• Cations migrate to the cathode.
• Anions migrate to the anode.

Redox Chemistry

Commercial Cells (Batteries)

• Devices which store chemical energy, which can be transformed by oxidation-reduction reactions into electrical energy.

• Three major types of cells:

  1. Primary ($1^o$) cells: May only be discharged (e.g., Dry Cell).

  2. Secondary ($2^o$) cells: May be recharged many times (e.g., Lead-acid Cell).

  3. Fuel cells: Continuous voltage generated by a constant flow of reactants (e.g., $H<em>2(g)$ and $O</em>2(g)$).

The Leclanche or Dry Cell (Primary Cell)

• Anode: $Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-$
• Cathode: $2MnO<em>2(s) + 2NH</em>4^+(aq) + 2e^- \rightarrow Mn<em>2O</em>3(s) + 2NH<em>3(g) + H</em>2O(l)$
• Overall: $2MnO<em>2(s) + 2NH</em>4^+(aq) + Zn(s) \rightarrow Mn<em>2O</em>3(s) + 2NH<em>3(g) + H</em>2O(l) + Zn^{2+}(aq)$

The Lead Accumulator (Secondary Cell)

• Anode: $Pb(s) + SO<em>4^{2-}(aq) \rightarrow PbSO</em>4(s) + 2e^-$
• Cathode: $PbO<em>2(s) + SO</em>4^{2-}(aq) + 4H^+(aq) + 2e^- \rightarrow PbSO<em>4(s) + 2H</em>2O(l)$
• Overall: $Pb(s) + 2SO<em>4^{2-} + PbO</em>2(s) + 4H^+(aq) \rightarrow 2PbSO<em>4(s) + 2H</em>2O(l)$

The Fuel Cell (Continuous Cell)

• (For an acid electrolyte)

• Anode: $H_2(g) \rightarrow 2H^+(aq) + 2e^-$

• Overall: $2H<em>2(g) + O</em>2(g) \rightarrow 2H_2O(l)$

Corrosion Chemistry

• Corrosion is an oxidation-reduction process that involves a metal and oxygen and possibly water.
• In many metals the oxide forms a layer on the metal surface that prevents the penetration of more oxygen and therefore protects the metal from further corrosion. Two examples are shown at right. Other examples are Zinc, Lead and Tin.

4 Al + 3$O<em>2$ OR 2 $Al</em>2O_3$

2 Cu + $O_2$ 2 CuO

The Corrosion of Iron

• Iron corrosion requires the presence of both water and oxygen.

Oxidation

• Fe (s) -> $Fe^{2+}$ (aq) + 2$e^-$

Reduction

*O2(g) + 2H2O (l) + 4e -> 4 OH(aq)

Precipitation

*Fe^{2+} (aq) + 2 OH(aq) -> Fe(OH)_2(s)

*Further oxidation of the resultant iron (II) hydroxide then occurs.

*4 Fe(OH) 2 (s) + 2 $H<em>2O$ (1) + O2(g) -> 4 FeO(OH).$$H_2O(s) (brown RUST)

Prevention of Rusting

Physical Means

*Covering the metal surface with oil or grease. eg: bicycle chains.
*Painting the metal surface. eg: motor vehicles.
*Coating the iron with a metal that forms a protective oxide. eg: Chromium handle bars on bicycles etc..

Chemical Means

*Anodic protection using a sacrificial zinc block that forms a positive potential at which oxidation occurs.
*Cathodic protection where an im- pressed charge results in the iron structure becoming the cathode of a circuit. (ie: oxidation cannot occur).
*Alloying iron with chromium, vana- dium or molybdenum to make stainless steels.
*Coating the metal with a polymeric coating.
*Forming insoluble iron compounds on the surface of the metal. eg: treatment with phosphoric acid forms iron (III) phosphate as an inert layer.