Oxidation & ReductionElectrochemistry BLB 11th Chapters 4, 20
Chapter Summary • Oxidation and Reduction (redox) – introduced in chapter 4 • Oxidation Numbers • Electron-transfer • Balancing redox reaction • Electrochemical cells • Corrosion • Electrolysis
20.1, 4.4 Oxidation-Reduction Reactions • Oxidation • Loss of electrons • Increase in oxidation number • Gain of oxygen or loss of hydrogen • Reduction • Gain of electrons • Decrease in oxidation number • Loss of oxygen or gain of hydrogen Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) • Oxidizing agent or oxidant – reactant that contains the element being reduced; is itself reduced • Reducing agent or reductant – reactant that contains the element being oxidized; is itself oxidized
Oxidation Numbers (p. 132) Assign according to the following order: • Atoms zero (since neutral) • Ions equal to charge of the ion • Nonmetals • O −2 • H +1 (when bonded to other nonmetals) −1 (when bonded to metals) • F −1 • X −1 except when combined with oxygen Sum of the oxidation numbers equals zero or the charge of the polyatomic ion.
O2 CH4 NO3¯ CH3OH Cr2O72- CH2O Cu2+ OCl¯ Oxidation numbers practice
Redox Reactions • Combustion, corrosion, metal production, bleaching, digestion, electrolysis • Metal oxidation • Activity Series (Table 4.5, p. 136) • Some metals are more easily oxidized and form compounds than other metals. • Displacement reaction – metal or metal ion is replaced through oxidation A + BX → AX + B
20.2 Balancing Redox Reactions • Goal: Balance both the atoms and the electrons • Examples: Al(s) + Zn2+(aq) → Al3+(aq) + Zn(s) MnO4¯(aq) + Cl¯(aq) → Mn2+(aq) + Cl2(g)
The Rules (p. 830-1) In acidic solution: • Divide equation into two half-reactions (ox and red). • Balance all elements but H and O. • Balance O by adding H2O. • Balance H by addingH+. • Balance charge by adding electrons (e-). • Cancel out electrons by integer multiplication. • Add half reactions & cancel out. • Check balance of elements and charge.
The Rules (p. 833) In basic solution: Proceed as for acidic solution through step 7. • Add OH¯ to neutralize the H+. (H+ + OH¯ → H2O) • Cancel out H2O. • Check balance of elements and charge.
20.3 Voltaic Cells • A spontaneous redox reaction can perform electrical work. • The half-reactions must be placed in separate containers, but connected externally. • This creates a potential for electrons to flow. • Reactant metal is the most reactive; product metal the least. Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) Line notation: Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s)
20.3 Voltaic Cell Net reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Movement of Electrons Zn(s) → Zn2+(aq) + 2 e¯ Cu2+(aq) + 2 e¯ → Cu(s) e¯ Net reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
20.4 Cell Potentials Under Standard Conditions • EMF – electromotive force – the potential energy difference between the two electrodes of a voltaic cell; Ecell; measured in volts • E°cell – standard cell potential (or standard emf) • For the Zn/Cu cell, E°cell = 1.10 V • electrical work = Coulombs x volts J = C x V
Standard Reduction (Half-cell) Potentials • E° - potential of each half-cell • E°cell = E°cell(cathode) -E°cell(anode) • For a product-favored reaction: • ΔG° < 0 • E°cell > 0 • Measured against standard hydrogen electrode (SHE); assigned E° = 0 V.
Problem Voltaic cell with: Al(s) in Al(NO3)3(aq) on one side and a SHE on the other. Sketch the cell, determine the balance equation, and calculate the cell potential.
Problem Voltaic cell with: Pb(s) in Pb(NO3)2(aq) on one side and a Pt(s) electrode in NaCl(aq) with Cl2 bubbled around the electrode on the other. Sketch the cell, determine the balance equation, and calculate the cell potential.
20.5 Free Energy and Redox Reactions • ΔG° < 0 • E°cell > 0 ΔG° for previous problems
20.6 Cell Potentials Under Nonstandard Conditions • Concentrations change as a cell runs. • When E = 0, the cell is dead and reaches equilibrium. • Nernst equation allows us to calculate E under nonstandard conditions:
Concentration Cells • A cell potential can be created by using same half-cell materials, but in different concentrations. Problem 69
Cell EMF and Equilibrium • When E = 0, no net change in flow of electrons and cell reaches equilibrium. K of previous problems
20.7 Batteries and Fuel Cells Batteries • self-contained electrochemical power source • More cells produce higher potentials • Primary – non-rechargeable (anode/cathode) • Alkaline: Zn in KOH/MnO2 • Secondary – rechargeable (anode/cathode) • Lead-acid: Pb/PbO2 in H2SO4 • nicad: Cd/[NiO(OH)] • NiMH: ZrNi2/[NiO(OH)] • Li-ion: C(s,graphite)/LiCoO2
Hydrogen Fuel Cells • Convert chemical energy directly into electricity • Fuel and oxidant supplied externally continuously • Products are only electricity and water cathode: O2(g) + 4 H+(aq) + 4 e¯ → 2 H2O(l) anode: 2 H2(g) → 4 H+(aq) + 4 e¯ overall: 2 H2(g) + O2(g) → 2 H2O(l)
20.8 Corrosion • RUST! • Anode: M(s) → Mn+(aq) + n e¯ • Cathode: O2(g) + 4 H+(aq) + 4 e¯ → 2 H2O(l) or: O2(g) + 2 H2O(l) + 4 e¯ → 4 OH¯ (aq)
Preventing Corrosion • Anionic inhibition • painting • oxide formation • coating • Cathodic inhibition • sacrificial anode – attach a metal (like Mg) more easily oxidized • galvanizing steel – coating with zinc
20.9 Electrolysis • Electrical energy chemical change
Hall-Héroult Process for Al Production C(s) + 2 O2-(l) → CO2(g) + 4 e¯ 3 e¯ + Al3+(l) → Al(l)