2 14-1 Redox Chemistry & Electricity Oxidation: a loss of electrons to an oxidizing agentReduction: a gain of electrons from a reducing agentReduction-Oxidation reaction (redox reaction)ex:Half-reactions:re: Fe3+ + e- Fe2+ox: V2+ V3+ + e-
3 1. Chemistry & Electricity Electrochemistry: the study of the interchange of chemical & electrical energy.Electric charge (q) is measured in coulombs(C).The magnitude of the charge of a single electron (or proton) is 1.602×10－19 C. A mole of electrons therefore has a charge of (1.602×10－19 C)(6.022×1023 /mol)= 9.649×104 C/mol, which is called the Faraday constant, F.Example at p.310
4 2. Electric current is proportional to the rate of a redox reaction I (ampere; A) = electric current= a flow of 1 coulomb per second = 1C/sExample at p. 310:Sn4+ + 2e- Sn2+at a constant rate of 4.24 mmole/h.How much current flows into the solution?
5 3. Voltage & Electrical Work The difference in electric potential between two points measures the work that is needed (or can be done) when electrons move from one point to another.wireqIEhoseH2OVH2OPH2O
6 Ask yourself at p.312Consider the redox reaction
7 14.2 Galvanic CellsChemical reaction spontaneously occurs to produce electrical energy.Ex: lead storage batteryWhen the oxidizing agent & reducing agent are physically separated, e transfer through an external wire. generates electricity.
8 A cell in action Electrodes: the redox rxn occur anode: oxidation occurcathode: reduction occurSalt bridge: connect two solns.External wire
9 Cell representation: Line Notation Example： Interpreting Line Diagrams of CellsFigure 14-4 Another galvanic cell.
10 14-3 Standard Potentials Cell potential ( Ecell) The voltage difference between the electrodes. electromotive force (emf)can be measured by voltmeter.emf of a cell depends onThe nature of the electrodes & [ions]Temp.
11 14-3 Standard Potentials S.H.E. (standard hydrogen electrode ) It is impossible to measure Ecell of a half-rxn directly, need a reference rxn.standard hydrogen electrode:
12 The standard reduction potential (E0) for each half-cell is measured by an experiment shown in idealized form in Fig.14-6.
15 Standard Reduction Potentials for reaction rxn:
16 Formal potential AgCl (s) + e- Ag (s) + Cl- 0.222 V0.197 V in saturated KCl (formal potentional)E0 = 0.222VS.H.E.║ Cl- (aq, 1M) | AgCl (s) | Ag(s)E0’ (formal potential) = V (in saturated KCl)S.H.E.║ KCl (aq, saturated) | AgCl (s) | Ag(s)
17 Formal Potential with H+A- E°≠1.61V Formal potential: (E°’) Ex: Ce4+ + e- Ce3+ E°=1.6Vwith H+A E°≠1.61VFormal potential: (E°’)The potential for a cell containing a [reagent] ≠1M.Ex: Ce4+/Ce3+ in 1M HCl E°’=1.28V
18 14-4 The Nernst EquationThe net driving force for a reaction is expressed by the Nernst eqn.Nernst Eqn for a Half-ReactionwhereE is the reduction potential at the specified concentrationsn: the number of electrons involved in the half-reactionR: gas constant ( V coul deg-1mol-1)T: absolute temperatureF: Faraday constant (96,487 coul eq-1) at 25°C RT/F=
19 Nernst equation for a half-reaction at 25ºC E = E0 when [A] = [B] = 1MQ (Reaction quotient ) =1 E = E0Where, Q = [B]b / [A]a
20 [C] & Ecell standard conditions: [C]=1M what if [C]≠1M? (ex) [Al3+]=2.0M, [Mn2+]=1.0M Ecell<0.48V[Al3+]=1.0M, [Mn2+]=3.0M Ecell>0.48V
21 Dependence of potential on pH Many redox reactions involved protons, and their potentials are influenced greatly by pH.
22 Nernst Equation for a Complete Reaction 1. Write reduction half-reactions for both half-cells and find E0 for each in Appendix C.2. Write Nernst equation for the half-reaction in the right half-cell.3. Write Nernst equation for the half-reaction in the left half-cell.4. Fine the net cell voltage by subtraction: E=E＋－ E－.5. To write a balanced net cell reaction.P.321
23 Nernst Equation for a complete reaction Example at p. 321Rxn: 2Ag+ (aq) + Cd (s) Ag (s) + Cd 2+ (aq)2Ag+ + 2e- Ag (s) E0+ = 0.799Cd e- Cd (s) E0- =
24 Electrons Flow Toward More Positive Potential Electrons always flow from left to right in a diagram like Figure 14-7.
27 At equilibriumE = 0 and Q = KE0 > 0 K > 1,E0 < 0, K < 1
28 Ex:One beaker contains a solution of M KMnO4, M MnSO4, and M H2SO4; and a second beaker contains M FeSO4 and M Fe2 (SO4)3. The 2 beakers are connected by a salt bridge and Pt electrodes are placed one in each. The electrodes are connected via a wire with a voltmeter in between.What would be the potential of each half-cell (a) before reaction and (b) after reaction?What would be the measured cell voltage (c) at the start of the reaction and (d) after the reaction reaches eq.?Assume H2SO4 to be completely ionized and equal volumes in each beaker.