Chapter 1: Redox Equilibrium

1.1 Oxidation and Reduction

• A redox reaction is a chemical reaction where oxidation and reduction occur simultaneously. (p.4)
• Oxygen: oxidation = gain of oxygen; reduction = loss of oxygen. (p.4)
• Example: 2Mg + CO₂ → 2MgO + C. Mg is oxidised (reducing agent); CO₂ is reduced (oxidising agent). (p.5)
• Hydrogen: oxidation = loss of hydrogen; reduction = gain of hydrogen. (p.5)
• Example: H₂S + Cl₂ → S + 2HCl. H₂S is oxidised; Cl₂ is reduced. (p.6)
• Electrons: oxidation = loss of electrons; reduction = gain of electrons. (p.6)
• Reducing agent is the electron donor; oxidising agent is the electron acceptor. (p.7)
• Transfer of electrons at a distance: electrons flow through the connecting wires from the reducing agent (negative terminal) to the oxidising agent (positive terminal); the galvanometer needle deflects. (p.8)
• Sulphuric acid separates the two solutions and acts as the electrolyte. (p.8)
• Oxidation half equation: 2I⁻ → I₂ + 2e⁻ (colourless to brown). (p.9)
• Reduction half equation: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O (purple to colourless). (p.9)
• Ionic equation: 10I⁻ + 2MnO₄⁻ + 16H⁺ → 5I₂ + 2Mn²⁺ + 8H₂O (electrons in both half equations must be equal before adding). (p.10)
• Common oxidising agents: acidified KMnO₄, acidified K₂Cr₂O₇, Cl₂, Br₂, FeCl₃, acidified H₂O₂. (p.10)
• Common reducing agents: KI, KBr, FeSO₄, reactive metals (e.g. Zn), SO₂. (p.10)

Oxidation Number Rules

• The oxidation number of all elements is zero. (p.11)
• The oxidation number of a monoatomic ion equals its charge. (p.11)
• Group 1, Group 2 and Group 13 elements in compounds are +1, +2 and +3 respectively. (p.11)
• The sum of oxidation numbers in a compound is zero. (p.11)
• The sum of oxidation numbers in a polyatomic ion equals the charge of the ion. (p.11)
• Group 17 in a compound is normally −1 (fluorine is always −1; Cl, Br, I are positive when bonded to oxygen). (p.12)
• Hydrogen is normally +1, except in metal hydrides (−1). (p.12)
• Oxygen is normally −2, except in peroxides (−1). (p.12)
• Oxidation = oxidation number increases; reduction = oxidation number decreases. (p.13)
• An oxidising agent oxidises other substances and is itself reduced; a reducing agent reduces other substances and is itself oxidised. (p.13)
• IUPAC naming: Roman numerals show the oxidation number of metals with more than one oxidation number, e.g. iron(II) nitrate, Fe(NO₃)₂; potassium manganate(VII), KMnO₄; potassium dichromate(VI), K₂Cr₂O₇. (p.14)
• Fe²⁺ can be converted to Fe³⁺ using bromine water; Fe³⁺ to Fe²⁺ using zinc powder. Presence tested with NaOH, NH₃, K₄[Fe(CN)₆], K₃[Fe(CN)₆] or KSCN. (p.15)
• Fe²⁺ → Fe³⁺ is oxidation (loses an electron, +2 to +3); Fe³⁺ → Fe²⁺ is reduction. (p.16)
• A more electropositive metal displaces a less electropositive metal from its salt solution. (p.16)
• Zn + Cu(NO₃)₂: blue solution becomes paler, brown solid (Cu) deposited, Zn plate becomes thinner; Cu²⁺ is reduced, Zn is oxidised. (p.16)
• Halogen oxidising strength decreases down Group 17 (F₂ strongest); halide ion reducing strength increases down the group. (p.19)
• Cl₂ + 2KBr → Br₂ + 2KCl: colourless KBr turns brown; ionic equation 2Br⁻ + Cl₂ → Br₂ + 2Cl⁻. (p.19)
• 1,1,1-trichloroethane is used to identify the halogen formed by its colour in the organic layer. (p.20)

1.2 Standard Electrode Potential

• Electrode potential is the potential difference produced when an equilibrium is established between metal M and its aqueous Mⁿ⁺ ions in a half-cell. (p.22)
• Standard conditions: ion concentration 1.0 mol dm⁻³; gas pressure 1 atm (101 kPa); temperature 25 °C (298 K); platinum as inert electrode when the half-cell is not a metal. (p.22)
• The standard hydrogen electrode is the reference electrode: Pt in 1.0 mol dm⁻³ H⁺ with H₂ at 1 atm. 2H⁺ + 2e⁻ ⇄ H₂, E° = 0.00 V. (p.23)
• Zn²⁺ + 2e⁻ ⇄ Zn, E° = −0.76 V (zinc is the negative terminal against the SHE). (p.23)
• Table 1.6 (E°, V): Li⁺/Li −3.04; K⁺/K −2.92; Ca²⁺/Ca −2.87; Na⁺/Na −2.71; Mg²⁺/Mg −2.38; Al³⁺/Al −1.66; Zn²⁺/Zn −0.76; Fe²⁺/Fe −0.44; Ni²⁺/Ni −0.25; Sn²⁺/Sn −0.14; Pb²⁺/Pb −0.13; H⁺/H₂ 0.00; Cu²⁺/Cu +0.34; I₂/I⁻ +0.54; Fe³⁺/Fe²⁺ +0.77; Ag⁺/Ag +0.80; Br₂/Br⁻ +1.07; Cr₂O₇²⁻/Cr³⁺ +1.33; Cl₂/Cl⁻ +1.36; MnO₄⁻/Mn²⁺ +1.52; H₂O₂ +1.77; F₂/F⁻ +2.87. (p.24)
• More negative E°: stronger reducing agent (easier to oxidise). More positive E°: stronger oxidising agent (easier to reduce). (p.24)
• Cu can displace Ag from AgNO₃ because Cu is a stronger reducing agent: Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag. (p.25)
• Cu cannot displace Mg from Mg(NO₃)₂ because Cu is a weaker reducing agent; no reaction occurs. (p.26)

1.3 Voltaic Cell

• A voltaic (galvanic) cell converts chemical energy to electrical energy through a redox reaction. (p.27)
• Negative terminal (anode): oxidation. Positive terminal (cathode): reduction. Electrons flow from negative to positive; current flows from positive to negative. (p.27)
• Daniell cell: Zn anode Zn → Zn²⁺ + 2e⁻; Cu cathode Cu²⁺ + 2e⁻ → Cu; overall Zn + Cu²⁺ → Zn²⁺ + Cu. (p.28)
• Cell notation: Electrode(s) | Electrolyte(aq) || Electrolyte(aq) | Electrode(s), with anode on the left and cathode on the right. (p.28)
• Daniell cell notation: Zn(s) | Zn²⁺(aq, 1.0 mol dm⁻³) || Cu²⁺(aq, 1.0 mol dm⁻³) | Cu(s). (p.29)
• E°_cell = E°(cathode) − E°(anode) = E°(positive terminal) − E°(negative terminal); Daniell cell = +0.34 − (−0.76) = +1.10 V. (p.29)

1.4 Electrolytic Cell

• Electrolytes conduct electricity in molten state or aqueous solution and undergo chemical changes; non-electrolytes cannot conduct in any state. (p.31)
• Electrolytes conduct because of freely moving ions; solid ionic compounds cannot conduct because ions are fixed in the lattice. (p.32)
• Conductors (metals, graphite) conduct by free electrons without chemical change, and conductivity decreases as temperature rises; electrolytes conduct by ions and conductivity increases with temperature. (p.32)
• Electrolysis is the decomposition of a compound in molten state or aqueous solution into its elements by passing electricity through it. Anions move to the anode, cations to the cathode. (p.34)
• Molten PbBr₂: anode 2Br⁻ → Br₂ + 2e⁻ (brown gas, oxidation); cathode Pb²⁺ + 2e⁻ → Pb (grey solid, reduction). (p.35)
• Three factors affecting product selection in aqueous electrolysis: E° value, concentration of ions, type of electrode. (p.37)
• Anode: the anion with a more negative/less positive E° is discharged. Cathode: the cation with a more positive/less negative E° is discharged. (p.37)
• If halide ions are present in higher concentration, they are discharged at the anode even if their E° is more positive. (p.37)
• Active electrodes (Cu, Ag) at the anode: the metal atoms ionise instead of anions being discharged. (p.37)
• Water dissociates: H₂O ⇄ H⁺ + OH⁻. (p.37)
• CuSO₄ solution with carbon electrodes: anode releases O₂ (colourless gas); cathode deposits brown copper. (p.39)
• Anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻. Cathode: Cu²⁺ + 2e⁻ → Cu. (p.40)
• Concentrated (1.0 mol dm⁻³) HCl gives Cl₂ at the anode; very dilute (0.0001 mol dm⁻³) HCl gives O₂. Chlorine bleaches damp blue litmus paper. (p.41)
• Type of electrode: with a copper anode in CuSO₄, Cu dissolves instead of OH⁻ being discharged. (p.42)
• Voltaic cell: chemical to electrical energy; two different metals; anode is negative. Electrolytic cell: electrical to chemical energy; electrodes may be the same; anode is connected to the positive terminal of the battery. (p.43)
• Electroplating: object = cathode, plating metal = anode, electrolyte contains ions of the plating metal. Anode Cu → Cu²⁺ + 2e⁻; cathode Cu²⁺ + 2e⁻ → Cu. (p.44)
• The electrolyte colour stays the same in electroplating because the Cu²⁺ concentration remains constant. (p.44)
• Purification of copper: impure Cu anode, pure Cu cathode, copper salt electrolyte. Anode becomes thinner and impurities collect below it; cathode becomes thicker. (p.46)

1.5 Extraction of Metal from Its Ore

• Very reactive metals (K, Na, Mg, Al) are extracted by electrolysis; moderately reactive metals (Zn, Fe, Sn, Pb) by reduction with carbon; Cu and Hg by heating in air; Ag and Au exist as elements. (p.48)
• Aluminium: bauxite is purified to Al₂O₃, then melted with cryolite, Na₃AlF₆, to lower the melting point from 2000 °C. (p.48)
• Anode: 2O²⁻ → O₂ + 4e⁻. Cathode: Al³⁺ + 3e⁻ → Al. Molten aluminium sinks and is drained out. (p.49)
• Environmental effects: huge electricity use, CO₂ formed, and toxic red sludge from bauxite purification; recycling aluminium reduces pollution. (p.49)
• Blast furnace: C + O₂ → CO₂ (exothermic, up to 1600 °C); C + CO₂ → 2CO; Fe₂O₃ + 3CO → 2Fe + 3CO₂. (p.50)
• CaCO₃ → CaO + CO₂; CaO + SiO₂ → CaSiO₃ (slag), which removes acidic impurities. (p.50)
• Thermite reaction: Fe₂O₃ + 2Al → Al₂O₃ + 2Fe(l); used to weld railway tracks. (p.51)

1.6 Rusting

• Corrosion of metal is a redox reaction where the metal is oxidised spontaneously: M → Mⁿ⁺ + ne⁻. The more electropositive the metal, the easier it corrodes. (p.52)
• Rusting needs water and oxygen; iron is the reducing agent, oxygen the oxidising agent. (p.53)
• Anode (low oxygen, middle of droplet): Fe → Fe²⁺ + 2e⁻. Cathode (high oxygen): O₂ + 2H₂O + 4e⁻ → 4OH⁻. (p.53)
• Fe²⁺ + 2OH⁻ → Fe(OH)₂, which is further oxidised to hydrated iron(III) oxide, Fe₂O₃.xH₂O (rust). (p.54)
• Rust is brittle and permeable so rusting continues; acids and salts speed up rusting by improving conductivity. (p.54)
• Contact with a more electropositive metal (Mg, Al, Zn) slows rusting; contact with a less electropositive metal (Sn, Pb, Cu) speeds it up. (p.55)
• Prevention: protective surface (paint, grease, plastic); sacrificial protection with Mg or Zn; galvanising (zinc), tin plating, chrome plating. (p.56)
• Alloying: stainless steel is iron alloyed with carbon, chromium and nickel, forming a protective oxide layer. (p.57)