Daniel L. Reger Scott R. Goode David W. Ball www.cengage.com/chemistry/reger Chapter 18 Electrochemistry Oxidation and Reduction • Oxidation is the loss of electrons by a chemical process. • When sodium forms a compound, Na+ is formed. Sodium is oxidized. • Reduction is the gain of electrons by a chemical process. • When Cl- ions are formed from elemental chlorine, chlorine is reduced. Oxidation-Reduction (“Redox”) • An oxidation-reduction reaction, or redox reaction, is one in which electrons are transferred from one species to another. • In every redox reaction, at least one species is oxidized and at least one species is reduced. • 2Na(s) + Cl2(g) → 2NaCl(s) is a redox reaction because Na is oxidized and Cl is reduced. Oxidizing and Reducing Agents • An oxidizing agent is the reactant that accepts electrons, causing an oxidation to occur. • The oxidizing agent is reduced. • A reducing agent is the reactant that supplies electrons, causing a reduction to occur. • The reducing agent is oxidized. • In the reaction of sodium with chlorine, Na is the reducing agent and Cl2 is the oxidizing agent. Half-reactions • In a half-reaction, either the oxidation or reduction part of a redox reaction is given, showing the electrons explicitly. • Half-reactions emphasize the transfer of electrons in a redox reaction. • For 2Na(s) + Cl2(g) → 2NaCl(s) : • Na → Na+ + 1e• Cl2 + 2e- → 2Cl- oxidation half-reaction reduction half-reaction Oxidation States • The oxidation state is the charge on the monatomic ion, or the charge on an atom when the shared electrons are assigned to the more electronegative atom. • Electron pairs shared by atoms of the same element are divided equally. • In CaCl2, an ionic compound: • calcium has an oxidation state of +2. • chlorine has an oxidation state of -1. Review • Assign oxidation numbers to each atom in the following substances. (a) PF3 (b) CO (c) NH4Cl Review • Balance the following equation in acid solution: • Cr2O72- + C2H5OH → Cr3+ + CO2 Review • Balance the following equation in basic solution: • Zn + ClO- → Zn(OH)42- + Cl- Voltaic Cells • A voltaic cell (also known as a galvanic cell) is an apparatus that produces electrical energy directly from a redox reaction. • All voltaic cells depend on redox reactions. Voltaic Cell - Diagram Voltaic Cells • The oxidation half-cell contains the reaction Zn(s) → Zn2+(aq) + 2e-. • The reduction half-cell contains the reaction Cu2+(aq) + 2e- → Cu(s). • Since the half-cells are physically separated, electrons must travel from one side to another through a connecting wire. Voltaic Cell Conventions • Assume that oxidation occurs in the left half-cell and reduction occurs in the right half-cell. • Zn(s) → Zn2+(aq) + 2e- occurs in left. • Cu2+(aq) + 2e- → Cu(s) occurs in right. • A voltmeter is connected to the wire coming out of each metal electrode. • If the measured voltage is positive, the chemical reaction is spontaneous as written. Electrodes and Half-Cells • Redox reactions not involving metals can be used in half-cells by using an inert electrode, like gold, platinum, or carbon to provide electrical contact. Examples of half-reactions include: • two soluble ions • Fe3+(aq) + e- → Fe2+(aq) • gas-ion reactions • Cl2(g) + e- → 2Cl-(aq) • insoluble salts • AgCl(s) + e- → Ag(s) + Cl-(aq) Example of Inert-Electrode Half-Cell The Hydrogen Electrode Electrical Potential • Electromotive force or emf (E) is the electrical driving force that “pushes” electrons from the oxidation half-cell to the reduction half-cell. • Cell potential is the potential energy per unit charge that is characteristic of each half-cell reaction. It is measured in volts (V). • 1 volt = 1 joule/coulomb (1 V = 1 J/C) Standard Potentials • The overall potential of a cell depends on the concentrations of the species in the reaction. • Standard potential, Ecell, is the cell potential when each species in the reaction is present in its standard state. • Solids, liquids, and gases are in their pure state at 1 atm pressure. • Solutes have 1 M concentration. Potentials of Voltaic Cells • For the reaction Zn(s) + Cu2+(aq) → Cu(s) + Zn2+(aq) the Ecell is +1.10 V. • A positive cell potential means that the reaction is spontaneous in the forward direction. Additivity of Cell Potentials • Cell potentials are additive: Cu(s) + 2Ag+(1 M) → 2Ag(s) + Cu2+(1 M) Zn(s) + Cu2+(1 M) → Zn2+(1 M) + Cu(s) +0.46 V +1.10 V Zn(s) + 2Ag+(1 M) → 2Ag(s) + Zn2+(1 M) E°cell =+1.10 V + 0.46 V = +1.56 V Standard Reduction Potential • The standard reduction potential of a half-reaction is the potential of the reduction reaction relative to the standard hydrogen electrode as the oxidation. Standard Reduction Potentials Table • Half-reactions written as reductions, where electrons are reactants. • Standard reduction potential for 2H+(aq,1 M) + 2e- → H2(g, 1 atm) is set to 0.000 V. • If a reduction reaction is reversed to make oxidation, change sign on reduction potential. Ag+(aq, 1 M) + e- → Ag(s) E = 0.80 V Ag(s) → Ag+(aq, 1 M) + e- E =-0.80 V Standard Reduction Potentials Table Reduction Half-Reaction F2(g) + 2e- → 2F-(aq) Ag+(aq) + e- → Ag(s) Fe3+(aq) + e- → Fe2+(aq) Sn4+(aq) + 2e- → Sn2+(aq) 2H+(aq) + 2e- → H2(g) Co2+(aq) + 2e- → Co(s) Fe2+(aq) + 2e- → Fe(s) Zn2+(aq) + 2e- → Zn(s) Mg2+(aq) + 2e- → Mg(s) E (V) 2.87 0.80 0.77 0.15 0.00 -0.28 -0.44 -0.76 -2.37 Calculating Cell Potential • The standard potential of a cell reaction is given by Ecell = Ered + Eox • Both half-reactions must transfer the same number of electrons. • Multiplying the coefficients of a halfreaction to balance the electrons does NOT change the potential. Fe3+(aq) + e- → Fe2+(aq) E = 0.77 V 2Fe3+(aq) + 2e- → 2Fe2+(aq) E = 0.77 V Example: Using Standard Potentials • Calculate the potential of the cell reaction 2Fe3+ + 2I- → 2Fe2+ + I2 from the potentials in Table 18.1. Test Your Skill • Write the spontaneous cell reaction and calculate Ecell for the voltaic cell made up from the two half-reactions below. Sn4+(aq) + 2e- → Sn2+(aq) Mg2+(aq) + 2e- → Mg(s) E = 0.15 V E = -2.37 V Test Your Skill • Write the spontaneous cell reaction and calculate Ecell for the voltaic cell made up from the two half-reactions below. Sn4+(aq) + 2e- → Sn2+(aq) Mg2+(aq) + 2e- → Mg(s) E = 0.15 V E = -2.37 V • Answer: Mg(s) + Sn4+(aq) → Mg2+(aq) + Sn2+(aq) Ecell = 2.52 V Activity Series • An activity series arranges halfreactions in order of decreasing potential. Use of Activity Series • Species with large positive reduction potentials are oxidizing agents – they oxidize species below them in the activity series. • Species with large negative reduction potentials are reducing agents – they reduce species above them in the activity series. Activity Series • What happens when Fe is added to 1 M solutions of Zn(NO3)2 and Co(NO3)2? Co2+(aq) + 2e- → Co(s) Fe2+(aq) + 2e- → Fe(s) Zn2+(aq) + 2e- → Zn(s) E = -0.28 V E = -0.44 V E = -0.76 V • Fe reacts with Co2+ but not with Zn2+. Test Your Skill • From the data in Table 18.1, select a metal that reduces Ag+(aq) but does not reduce AgCl(s). Write the reaction and calculate its standard potential. Cell Potentials and DG • DG = -nFE where DG = free energy change; n = number of electrons transferred; F = Faraday constant, 96,485 C/mol e-; E = cell potential. Cell Potentials and DG • Under standard conditions, DG = -nFE. • This means that spontaneous reactions have positive cell potentials. Calculating DG • Calculate DG for the reaction below. AgCl(s) → Ag+(aq) + Cl-(aq) using the following information: AgCl(s) + e Ag(s) + Cl (aq) + Ag(s) Ag (aq) + e - - E° = 0.222 V E° = -0.80 V Test Your Skill • Use the data in Appendix H to calculate DG for the following reaction. 2Na(s) + 2H2O(l) → 2Na+(aq) + 2OH-(aq) + H2(g) Relation of E to Keq • DG° = -RT ln Keq • DG° = -nFE° • -nFE° = -RT ln Keq • E° = RT nF ln Keq = • At 298 K, E° = 2 .3 0 3 R T nF 0 .0 5 9 1 n log Keq log Keq Calculating Keq from E • Determine Keq for the following reaction. Fe(s) + Pb2+(aq) → Pb(s) + Fe2+(aq) Test Your Skill • Determine Keq for the reaction below at 25 C. 2Ag(s) + Ni2+(aq) → Ni(s) + 2Ag+(aq) The Nernst Equation • The Nernst equation is used to calculate cell potentials under non-standard conditions: E ce ll E ce ll 2 .3 0 3 R T nF lo g Q E ce ll 0 .0 5 9 1 lo g Q n where Q is the reaction quotient. The second expression is for 25 C only. The Nernst Equation • What is the cell potential for a cell composed of Zn, 0.500 M Zn(NO3)2, Cu, and 1.500 M Cu(NO3)2? Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) E = 1.10 V Test Your Skill • What is the cell potential for a cell composed of Cu, 1.250 M Cu(NO3)2, Fe, and 0.100 M FeSO4? Applications of Voltaic Cells • Measurement of species concentration: the pH meter. • The half-cell potential depends on the concentration of hydrogen ions in solution. • A pH electrode is an electrochemical cell with a known reference cell as a part. • A pH meter is simply a volt meter calibrated to display pH instead of volts. Applications: Batteries • All batteries are voltaic cells. • In a dry cell, a Zn case is the anode, which is in contact with a moist paste of MnO2, NH4Cl, and a carbon electrode. • Cell reaction: Zn(s) + 2NH4+(aq) + MnO2(s) → Zn2+(aq) + Mn2O3(s) + 2NH3(aq) + H2O(l) Dry Cell Alkaline Dry Cell • • • • NH4Cl is replaced with KOH or NaOH. Zn anode corrodes less. Voltage is more constant. Overall cell reaction: Zn(s) + MnO2(s) → ZnO(s) + Mn2O3(s) Alkaline Dry Cell Lead Storage Battery • Used for high current applications. • Overall reaction: Pb(s) + PbO2(s) + 4H+ + 2SO42- → 2PbSO4(s) + 2H2O(l) • No salt bridge needed, because oxidizing and reducing agents do not come into contact with each other. • Rechargeable as long as the PbSO4 product adheres to electrodes (not limitless). Lead Storage Battery Fuel Cells • A fuel cell is a voltaic cell in which reactants are supplied continuously, products are removed continuously, and electrical energy is produced from the chemical reaction. • The H2/O2 fuel cell uses the reaction 2H2(g) + O2(g) → 2H2O(l) and is used to produce electricity and water on space missions.