Chapter 6 OXIDATION & REDUCTION

REDOX oxidation Reduction loss of electrons gain of electrons (RIG) or (GER) (OIL) or (LEO)

Reduction Oxidation Oxidant + e- Product Product + e- Reductant ( Gain of Electrons) (oxidation number Decreases) Br + e- Br - Oxidation Product + e- Reductant ( Loss of Electrons) (oxidation number Increases) K K+ + e-

(electron transfer reactions) 2Mg (s) + O2 (g) 2MgO (s) 2Mg 2Mg2+ + 4e- Oxidation half-reaction (lose e-) O2 + 4e- 2O2- Reduction half-reaction (gain e-) 2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e- 2Mg + O2 2MgO

4.4

Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s) Zn Zn2+ + 2e- Zn is oxidized Zn is the reducing agent Cu2+ + 2e- Cu Cu2+ is reduced CuSO4 is the oxidizing agent 4.4

Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s) Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s) Cu Cu2+ + 2e- Cu is the reducing agent Cu is oxidized Ag+ + 1e- Ag Ag+ is reduced AgNO3 is the oxidizing agent

2FeCl2(aq) + Cl2(g) 2FeCl3(aq) The ionic equation 2Fe+2(aq) + 4Cl-(aq) + Cl2(g) 2Fe+3(aq) + 6Cl-(aq) Spectator ions are removed 2Fe+2(aq) + Cl2(g) 2Fe+3(aq) + 2Cl-(aq) oxidised reduced The half equations for the reaction 2 Fe+2 2Fe+3 + 2e- (half equation) Cl2 + 2e- 2Cl- (half equation) 2Fe2+ + Cl2 2Fe3+ + 2Cl- (redox equation)

Q2: Look at the following reactions, for each one write down the ionic equation. Decide if the equation is a redox reaction. If it is a redox reaction write down the two half equations and state what has been oxidised and what has been reduced. 2Mg(s) + O2(g) 2MgO(s) Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s) HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Oxidation; increase in the oxidation number 2Fe2+ + Cl2 2Fe3+ + 2Cl- Reduction; decrease in the oxidation number

The differences between neutralization and redox reactions Neutralisation Redox reactions Involves proton transfer An acid is a proton donor A base is a proton acceptor Involve electron transfer An oxidising agent is an electron acceptor A reducing agent is an electron donor Oxidation is the removal of electrons. Reduction is a gain of electrons.

Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. Free elements (uncombined state) have an oxidation number of zero. Na (s), Be, K(s), Pb (s), H2 (g), O2 (g) , Hg (l), P4 = 0 In monatomic ions, the oxidation number is equal to the charge on the ion. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 The oxidation number of oxygen is usually –2. In H2O2 and O22- it is –1. 4.4

HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4 The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. Group IA metals are +1, IIA metals are +2 and fluorine is always –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. HCO3- Oxidation numbers of all the elements in HCO3- ? O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4 4.4

Rules for Assigning Oxidation States

Figure 4.10 The oxidation numbers of elements in their compounds 4.4

IF7 F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 Na = +1 O = -2 O = -2 Oxidation numbers of all the elements in the following ? F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 Na = +1 O = -2 O = -2 K = +1 3x(-2) + 1 + ? = 0 7x(-2) + 2x(+1) + 2x(?) = 0 I = +5 Cr = +6 4.4

Example What is the oxidation numbers of Cr and N in K2Cr2O7 and HNO3 respectively? K2Cr2O7 HNO3 +1 - 2 +1 - 2 Cr in K2Cr2O7 2 (+1) + 2n + 7 (-2) = 0 n = +6 N in HNO3 (+1) + n + 3(-2) =0 n = +5

Q 3: Find the oxidation numbers of: Cr in CrO42- ii. Mn in Mn(s) iii. Fe in FeCl3 iv. P in P4O6 v. Br in Br2(l) CrO42- 4x(-2) + ? = -2 Cr = +6 O = -2 Fe = +3 3x(-1) + ? = 0 P = +3 6x(-2) + 4n = 0

Oxidising and reducing agents oxidising agent Reducing agent an element or compound that oxidises another element or compound an element or compound that reduces another element or compound It is reduced It is oxidised

Oxidising and reducing agents 2FeCl2(aq) + Cl2(g) 2FeCl3(aq) reducing agent oxidising agent NOTE when the oxidising and reducing agent is named, the whole compound is specified, not just the element that undergoes the change in the oxidation state.

Q4: In the following reactions state which is the oxidising agent and which is the reducing agent. (It will help to write down the half equations) 2Mg(s) + O2(g) 2MgO(s) Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s) Cl2(g) + 2KI(aq) 2KCl(aq) + I2(s)

Oxidising agents Reducing agents 1. Liberate iodine from aqueous potassium iodide solution O.A 2I-(aq) I2(aq) + 2e- colourless brown Turn starch-iodide paper blue Decolorize acidified potassium permanganate (VII) paper: R.A MnO-4(aq) + 8H+ + 5e- Mn+2 + 4H2O Purple colourless 2. Convert hydrogen sulphide to sulphur: S-2(aq) S(s) + 2e- Turn potassium dichromate (VI) paper green Cr2O7-2(aq) +14H+(aq) + 6e- 2Cr+3(aq +7H2O Orange Green 3. Convert Fe(II) salts to Fe(III): Fe+3 ions give a dark-blue precipitate with potassium hexacyanoferrate(II) (ferrocyanide), K4 Fe(CN)6 , or a red coloration with aqueous potassium thiocyanate solution, KSCN. 3. Convert Fe(III) salts to Fe(II): Fe+2 salts give a dark-blue precipitate with potassium hexacyano ferrate(III) (ferricyanide), K3Fe(CN)6

Q5: Classify the following processes as redox (electron transfer), precipitation (ion transfer) or neutralisation (proton transfer) reactions: 2AgNO3(aq) + K2 CrO4(aq) Ag2 CrO4(s) + 2KNO3(aq) 2Al(s) + Fe2 O3(s) Al2 O3(s) + 2Fe(s) 3Fe(s) + 4H2 O(g) Fe3 O4(s) + 4H2(g) Ba(OH)2(aq) + 2HNO3(aq) Ba(NO3)2(aq) + 2H2O(l) BaCl2(aq) + H2 SO4(aq) BaSO4(s) + 2HCl(aq) Precipitation reaction Redox reaction Redox reaction Neutralisation reactions Precipitation reaction

Redox Process in Industry Oxygen (air) as oxidising agent Carbon or carbon monoxide as reducing agent 1. Methane is converted to methanol by direct oxidation 200 ˚C/copper tube 2CH4(g) + O2(g) 2 CH3OH(l)

Redox Process in Industry 2. Sulphuric acid is manufactured by oxidising the sulphur dioxide burn S(s) + O2(g) SO2(g) 420 ○C / V2 O5 2SO2(g) + O2(g) 2SO3(g) Then the product is absorbed in 98% acid and diluted with water. SO3(g) + H2 O H2 SO4

Redox Process in Industry 3. Carbon and carbon monoxide are used to reduce oxide ores to the metal Example: Iron(III) oxide is reduced to iron in a blast furnace. Fe2O3 + 3CO 2Fe + 3CO2 Iron (III) oxide is reduced to iron in a blast furnace

Redox Process in Industry 4- Aluminium oxide requires a more powerful reducing agent than either carbon or carbon monoxide. Aluminium is obtained from its ore by electrolytic reduction. At cathode Al3+ + 3e- Al Electrolytic reduction of Aluminium