Electrochemical cells = electronic conductor If two different + surrounding electrolytes are used: electrolyte electrode compartment Galvanic cell: electrochemical cell in which electricity is produced as a result of a spontaneous reaction (e.g., batteries, fuel cells, electric fish!) Electrolytic cell: electrochemical cell in which a non-spontaneous reaction is driven by an external source of current Nils Walter: Chem 260 Reactions at electrodes: Half-reactions Redox reactions: Reactions in which electrons are transferred from one species to another +II -II 00+IV -II → E.g., CuS(s) + O2(g) Cu(s) + SO2(g) reduced oxidized Any redox reactions can be expressed as the difference between two reduction half-reactions in which e- are taken up Reduction of Cu2+: Cu2+(aq) + 2e- → Cu(s) Reduction of Zn2+: Zn2+(aq) + 2e- → Zn(s) Difference: Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) - + - → 2+ More complex: MnO4 (aq) + 8H + 5e Mn (aq) + 4H2O(l) Half-reactions are only a formal way of writing a redox reaction Nils Walter: Chem 260 Carrying the concept further Reduction of Cu2+: Cu2+(aq) + 2e- → Cu(s) In general: redox couple Ox/Red, half-reaction Ox + νe- → Red Any reaction can be expressed in redox half-reactions: + - → 2 H (aq) + 2e H2(g, pf) + - → 2 H (aq) + 2e H2(g, pi) → Expansion of gas: H2(g, pi) H2(g, pf) AgCl(s) + e- → Ag(s) + Cl-(aq) Ag+(aq) + e- → Ag(s) Dissolution of a sparingly soluble salt: AgCl(s) → Ag+(aq) + Cl-(aq) − 1 1 Reaction quotients: Q = a − ≈ [Cl ] Q = ≈ Cl + a + [Ag ] Ag Nils Walter: Chem 260 Reactions at electrodes Galvanic cell: Electrolytic cell: Red → Ox + νe- 1 2 Red → Ox + νe- ν - → 1 2 Ox1 + e Red2 ν - → Ox1 + e Red2 Half-reactions Nils Walter: Chem 260 Types of electrodes I Gas electrode: Insoluble-salt electrode: Porous, insoluble Gas (e.g., H ) 2 salt (e.g., AgCl) solution metal solution metal - (e.g., H+) (e.g., Pt) (e.g., Cl ) (e.g., Ag) ||| ||| - = ||| ||| + Ag(s) AgCl(s) Cl (aq) Q a − Pt(s) H2(g) H (aq) Cl p(H ) - - + - = 2 AgCl(s) + e Ag(s) + Cl (aq) 2H + 2e H2(g) Q 2 + Nils Walter: Chem 260 aH Types of electrodes and how to put them together in a galvanic cell Redox electrode: Daniell cell: Left: Zn2+(aq) + 2e- → Zn(s) Right: Cu2+(aq) + 2e- → Cu(s) Pt(s)||| Fe2+(aq),Fe3+(aq) ||| ||| ||| ||| Zn(s) ZnSO4(aq) CuSO4(aq) Cu(s) 3+ - 2+ Fe + e Fe Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) a a 2+ a 2+ Q = Red = Fe Q = Zn Nils Walter: Chem 260 a a 3+ a 2+ Ox Fe Cu Cell reaction and potential Cathode (Right): Cu2+(aq) + 2e- → Cu(s) this way the reaction Anode (Left): Zn2+(aq) + 2e- → Zn(s) } becomes spontaneous Overall (R-L): Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) e- disappear Cell reaction: Difference of electrode half-reactions (Reduction at Cathode - Oxidation at Anode) Cell potential E: Potential difference between the electrodes ν - → Þ ν - Ox + e Red NA e transferred per mole of reaction Faraday constant = eNA = 96,485 Cmol-1 Þ ν × ν NA (-e) = - F charge transferred Maximum electrical work done in νNils ×Walter:∆ Chem 260 a galvanic cell: w’ = - F E = rG.