Electrochemical Cells

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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.

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Electrolytic Cells
CHEMISTRY LEVEL 4C (CHM415115) ELECTROLYTIC CELLS THEORY SUMMARY & REVISION QUESTIONS (CRITERION 5) ©JAK DENNY Tasmanian TCE Chemistry Revision Guides by Jak Denny are licensed under a Creative Commons Attribution-NonCommercial-NoDerivatives 4.0 International License. INDEX: PAGES • INTRODUCTORY THEORY 3 • COMPARING ENERGY CONVERSIONS 4 • APPLICATIONS OF ELECTROLYSIS 4 • ELECTROPLATING 5 • COMPARING CELLS 6 • THE ELECTROCHEMICAL SERIES 7 • PREDICTING ELECTROLYSIS PRODUCTS 8-9 • ELECTROLYSIS PREDICTION FLOWCHART 1 0 • ELECTROLYSIS PREDICTION EXAMPLES 11 • ELECTROLYSIS PREDICTION QUESTIONS 12 • INDUSTRIAL ELECTROLYTIC PROCESSES 13 • IMPORTANT ELECTRICAL THEORY 14 • FARADAY’S ELECTROLYSIS LAWS 15 • QUANTITATIVE ELECTROLYSIS 16 • CELLS IN SERIES 17 • ELECTROLYSIS QUESTIONS 18-20 • ELECTROLYSIS TEST QUESTIONS 21-22 • TEST ANSWERS 23 2 ©JAK CHEMISTRY LEVEL 4C (CHM415115) ELECTROLYTIC CELLS (CRITERION 5) INTRODUCTION: In our recent investigation of electrochemical cells, we encountered spontaneous redox reactions that release electrical energy such as takes place in the familiar situations of “batteries”. When a car battery is ‘flat’ and needs to be recharged, a power supply (‘battery charger’) is connected to the flat battery and chemical changes take place and it is subsequently able to operate again as a power supply. The ‘recharging’ process is non-spontaneous and requires an input of energy to occur. When a cell is such that an energy input is required to make a non-spontaneous redox reaction take place, we describe the cell as an ELECTROLYTIC CELL. The redox process occurring in an ELECTROLYTIC CELL is referred to as ELECTROLYSIS. For example, consider the spontaneous redox reaction associated with an electrochemical (fuel) cell; i.e. 2H2(g) + O2(g) → 2H2O(l) + ELECTRICAL ENERGY This chemical reaction RELEASES energy which can be used to power an electric motor, to drive a machine, appliance or car.
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