1. Chapter 7 Electrochemistry 7.1 Thermodynamic Properties of Electrolyte Solutions 7.1.1 Electrolyte Strong electrolyte Weak el ectrolyte Real electrolyte Potential electrolyte NaNO 3 z + ＝ 1 ｜ z - ｜＝ 1 1-1 ； BaSO 4 z + ＝ 2 ｜ z - ｜＝ 2 2-2 ； Na 2 SO 4 z + ＝ 1 ｜ z - ｜＝ 2 1-2 ； Ba(NO 3 ) 2 z + ＝ 2 ｜ z - ｜＝ 1 2-1 。

2. 7.1.2 Chemical Potential of Electrolyte and Ions  B = +  + +  

3. d T=0, d p=0, d n A =0  B = +  + + 

4. 7.1.3 Activity and Activity Coefficient ideal solution real solution

6. 7.1.4 Mean Activity of Ions and Mean Activity Coefficients

8. 7.1.5 The Debye - Hückel Limiting Law Ionic atmosphere

9. H2OH2O b<0.01 ～ 0.001mol·kg -1 I — Ionic Strength

10. 7.1.6 Ionic Strength I<0.01mol·kg -1

11. 7.2 Conductive Properties of Electrolyte Solutions 7.2.1 Conductance G  Conductance;unit Simens S ， 1S=1Ω -1 。  Resistivity ; Ω ·m.   Conductivity ; S·m -1.  =K (l/A) G K  Cell constant

12. 7.2.2 Molar Conductance Λ m unit S · m 2 · mol -1 。 Λ m (K 2 SO 4 ）＝ 0.02485 S·m 2 ·mol -1 Λ m ( K 2 SO 4 ）＝ 0.01243 S · m 2 · mol -1

13. 7.2.3 Concentration dependence of  and Λ m k/(S  m -1 ) H 2 SO 4 KOH KCl MgSO 4 CH 3 COOH 0145 10 c/(mol  dm -3 ) 20 40 60 80   c  m  c 400 300 200 100  m /(S  cm 2  mol -1 ) HCl NaOH AgNO 3 CH 3 COOH 0 0.5 1.0 1.5 c B =0 molar conductivity of infinite dilution

15. 7.2.4 Independent Migration of Ion 7.2.5. Electrolytic Equilibrium of Weak Electrolytes At equilibrium

16. a u =  u b u /b  ＝（ 1-α ）  u b/b  HOAc H + + OAc -

17. 7.3 Electrochemical system + + + + + + + + + + Metal + + + + + + + + + + + + + + + + + + + + + + + + + + + + + +

18. + + + + + + + + + + Metal 1Metal 2 Contact potential

19. Liquid-junction potential (diffusion potential) + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + ++ + + + + + + + + + + + + + + +

20. 7.4 Equilibrium electrochemistry 7.4.1 Reversible cell (1) Electrode reactions and cell reaction are reversible (2) I  0 (equilibrium)

21. 7.4.2 The Cell Potentials of Reversible Cell

22. 7.4.3 The Nernst Equation ---Standard Cell Potentials

24. 7.4.4 Standard Electrode Potential Standard Hydrogen Electrode ---SHE H + [a(H + ) =1 ] | H 2 (p  =100kPa) | Pt E  =0 SHE  electrode in question (reduction)

25. Table 11-1 25 ℃时某些电极的标准电极电势 (p  ＝ 100kPa) 电极电极反应（还原） E  /V K+KK+K K + + e - == K-2.924 Na +  Na Na + + e - == Na-2.7107 Mg 2+  Mg Mg 2+ + 2e - == Mg -2.375 Mn 2+  Mn Mn 2+ + 2e - == Mn -1.029 Zn 2+  Zn Zn 2+ + 2e - == Zn -0.7626 Fe 2+  Fe Fe 2+ + 2e - == Fe -0.409 Co 2+  Co Co 2+ + 2e - == Co -0.28 Ni 2+  Ni Ni 2+ + 2e - == Ni -0.23 Sn 2+  Sn Sn 2+ +2e - == Sn -0.1362 Pb 2+  Pb Pb 2+ +2e - == Pb -0.1261 H +  H 2  Pt H + +e - == 1/2H 2 -0.0000( 定义量 ) Cu 2+  Cu Cu 2+ +2e - == Cu +0.3402 Cu +  Cu Cu + +e - == Cu +0.522

26. Hg 2 2+  Hg Hg 2+ +2e - ==Hg+0.851 Ag +  Ag Ag + +e - == Ag+0.7991 OH -  O 2  Pt 1/2O 2 +H 2 O+2e - ==2OH - +0.401 H +  O 2  Pt O 2 +4H + + 2e - ==H 2 O+1.229 I -  I 2  Pt 1/2I 2 + e - == I - +0.535 Br -  Br 2  Pt 1/2Br 2 + e - ==Br - +1.065 Cl -  Cl 2  Pt 1/2Cl 2 + e - == Cl - +1.3586 I -  AgI  Ag AgI + e - ==Ag+I - -0.1517 Br -  AgBr  Ag AgBr + e - ==Ag+Br - +0.0715 Cl -  AgCl  Ag AgCl + e - ==Ag+Cl - +0.2225 Cl -  Hg 2 Cl 2  Hg Hg 2 Cl 2 + 2e - == 2Hg+2Cl - +0.2676 OH -  Ag 2 O  Ag Ag 2 O+2e - ==2Ag+2OH - +0.342 SO 4 2-  Hg 2 SO 4  Hg Hg 2 SO 4 +2e - == 2Hg+2SO 4 2- +0.6258 SO 4 2-  PbSO 4  Pb PbSO 4 + 2e - == Pb +SO 4 2- -0.356

27. Oxidation state + 2e -  Reduction state E MF = E (R, Reduction)- E (L, Reduction) Cl - - (a) ｜ Cl 2 ｜ Pt ： For example

28. 7.5 Application of E MF Measurements 7.5.1 Determination of thermodynamics quantities Δ r G m,Δ r S m andΔ r H m Δ r G m = - zFE MF Temperature coefficient of cell

29. 7.5.2 Determination of γ ±

30. 7.5.3 Determination of pH H + ｜ Q ， QH 2 ｜ Pt Q ［ a(Q) ］ +2H + ［ a(H + )]+2e -  QH 2 ［ a(QH 2 ) ］ 25 ℃， E ＝ (0.6997 - 0.05916pH) V Q HO - - OH QH 2 Q  QH 2 a ( Q)≈a(QH 2 ) Pt ｜ H 2 (p  ) ｜ solution(pH=x) ｜ KCl (a) ｜ Hg 2 Cl 2 ｜ Hg

31. 7.5.4 Determination of K  and K sp 7.5.5 Determination of reaction direction Δ r G m = - ZFE MF < 0

32. 7.6 kinetics of electrochemical system 7.6.1 Rate of electrochemical reaction M + + e - M  ca ca EcEc EaEa M a ca c M + +e M Cathode process υ c ； anode process υ a ；

33. v - Rate of electrochemical reaction mol  m -2  s -1 Current density j j=ZFυ j 0 :exchange current density

34. 7.6.2 Polarization and Overpotential { a}{ a} {  c,e } {  a,e } { c}{ c} { e}{ e}{  }     { a}{ a} { c}{ c} {j}{j} (a) electrolytic cell { }{ } { a}{ a} {  a,e } {  c,e } { c}{ c} { e}{ e}{  }     {j}{j} (b)chemical electric source { c}{ c} { a}{ a} { }{ } polarization curve Overpotential ： η a — anode overpotential η c — anode overpotential

35. (1). Diffusion overpotential Ag + c0c0 c'c'  Diffusion layer Ag M + + e - M (2). Electrochemical overpotential

36. 7.6.3 Electrolytic cell (- ） Pt ｜ H 2 ｜ OH - （ H 2 O ）｜ O 2 (p) ｜ Pt （ + ） H2H2 O2O2 H2OH2OH2OH2O PtPt PtPt anode(+)cathode(-) + — I Power supply

37. Pt A V R KOH 外电源 电阻 伏特计 电流计 + _ KOH VdVd  V  II Decomposition voltage Theory decomposition voltage Real decomposition voltage Δ  (real)=Δ  (theory) + (η a + ｜ η c ｜ ) + IR

38. 7.7 Power production and corrosion 7.7.1 Dry Cell Zn ｜ NH 4 Cl ｜ MnO 2 ｜ C Negative electrode ： Zn + 2NH 4 Cl  Zn(NH 3 ) 2 Cl 2 + 2H + + 2e - positive electrode ： 2MnO 2 + 2H + + 2e -  2MnOOH Cell reaction ： Zn + 2MnO 2 + 2NH 4 Cl  Zn(NH 3 ) 2 Cl 2 + 2MnOOH

39. 11.7.2 Storage Cell Pb ｜ H 2 SO 4 (ρ ＝ 1.28g  cm -3 ) ｜ PbO 2 Negative electrode ： Pb + H 2 SO 4  PbSO 4 + 2H + + 2e - positive electrode : PbO 2 + H 2 SO 4 + 2H + + 2e -  PbSO 4 + 2H 2 O Cell reaction ： PbO 2 + Pb + 2H 2 SO 4 2PbSO 4 + 2H 2 O

40. 11.7.3.Silver-zinc Cell Zn ｜ KOH(ω B ＝ 0.40) ｜ Ag 2 O ｜ Ag Negative electrode: 2Zn + 4OH -  2Zn(OH) 2 + 4e - positive electrode : Ag 2 O 2 + 2H 2 O + 4e -  2Ag + 4OH - Cell reaction ： 2Zn + Ag 2 O 2 + 2H 2 O 2Ag + 2Zn(OH) 2

41. 7.7.4. Fuel cell M ｜ H 2 (g) ｜ KOH ｜ O 2 (g) ｜ M

42. Efficiency of Chemical Electric Source

43. 7.7.5 Electrochemical corrosion

44. M+M+ 2H + H2H2 2e M M+M+ 2H + H2H2 2e M1M1 M2M2 Anode process ： Fe  Fe 2 ＋ +2e - Cathode process ： (i) 2H + +2e -  H 2 ↑ (ii) O 2 +4H + +4e -  2H 2 O (i) cell reaction ： Fe+2H ＋  Fe 2 ＋ +H 2 (ii) cell reaction ： Fe+(1/2)O 2 +2H +  Fe 2+ +H 2 O

45. {  )} S I {  c,e } {  a,e }