1. Chapter 7 Electrochemistry
Electrolysis
Electric energy Chemical energy
Galvanic cell

2. Definition of electrochemistry
A science studying the relationship between chemical energy and electrical energy and the rules of conversion of two energies.
II:
Electrochemistry is the study of solutions of electrolytes and of phenomena occurring at electrodes immersed in these solutions.

3. The application of electrochemistry
⒈ Electrolysis
hydrometallurgy of metal；
electrolytic preparation；
electroplating；
electrolytic oxidation.
⒉　Batteries
Primary batteries
Secondary Batteries
⒊　Electro-analysis
polarography，pH，electric conductivity, etc.

4. 7.1 Faraday’s law
Michael Faraday, a British chemist and physicist, studied the decomposition of solutions of salts, acids and bases by electric current.

5. Faraday's law of electrolysis states that the amount of any substance that is deposited or dissolved in electrolysis is directly proportional to the total passed electric charge.
It can be expressed as
F is Faraday constant.
F = Le = C·mol-1 = 1 Faraday.

6. coulometer
The silver coulometer and copper coulometer are commonly used ones.
If one mole of silver, g, is deposited on the cathode of a silver coulometer, we will know that the quantity of electricity passed through the coulometer is 1 coulomb.

7. 7.2 Transport numbers and mobilities
The definition of transport number t is that current carried by ion B divided by the sum of the current of all the ions in solution, which is also called the transference number of ion B.
nc and na are the amount of substance of positive ions migrate out from anodic region and that of negative ions migrate out from cathodic region respectively.

8. Let 4F be passed through the cell; t+=3t-
Before electrolysis
On electrolysis
After electrolysis

9. Mobility
The transport numbers of ions depend on the properties of ions and solvents, temperature, concentration, electric field strength and the like.
The mobility uB of an ion B is defined as its velocity in the direction of an electric field E of unit strength.

11. mobilities of hydrogen and hydroxyl ions
the mobility of hydrogen and hydroxyl ions in aqueous solution is abnormally high.
This is because both the H3O+ ion and hydroxyl ion are able to transfer a proton to a neighboring water molecule. This is explained by the fact that the H+ and OH- ions need not migrate through a protic solvent, but move through exchange of a proton between neighboring solvent molecules.
The unit of mobility is m2·V-1·s-1.

12. measurement of transport numbers by Hittorf method
The method of Hittorf is based on concentration changes in the anodic region and cathodic region in an electrolytic cell, caused by the passage of current through the electrolyte.

13. measurement of transport number by the moving boundary method
Suppose the boundary moves a distant x from AA’ to BB’ for the passage of Q coulombs. All the ions, H+, passed through the boundary AA’.
The amount of substances transported is then Q/F, of which t+Q/F are carried by the positive ion. If the volume between the boundaries AA’ and BB’ is V, and the concentration of HCl is c, then

14. 7.3 Conductance, conductivity and molar conductivity
7.3.1 The definition and the measurement
Conductance G in S
κ in equation is conductivity. In S m-1.
Molar conductivity (in S m2 mol-1)

15. Measurement of the conductivity

16. Application of conductivity measurement
Example The conductivity and molar conductivity of a saturated aqueous solution of silver chloride are 3.41×10-4S·m-1 and ×10-4S·m2·mol-1 respectively at 25℃. The conductivity of the water used to make the solution is 1.60×10-4S·m-1 at the same temperature. Calculate the solubility of silver chloride in water at 25℃.

17. Solution the conductivity of the silver chloride solution should be the sum of conductivities of silver chloride and water.
The solubility of silver chloride solution is then obtained

18. 7.3.2 The dependence of molar conductivity on the concentration
Fro strong electrolytes, Kohlrausch observed that m decreased with concentration according to the expression

20. conductivities and the concentrations of electrolytes

21. 7.3.3 Law of the independent migration of ions
Kohlrausch discovered relations between the values of for different electrolytes. For example
The difference in for pairs of salts having common ion is always approximately constant.

22. This behavior indicates that ions in an extremely dilute solution migrate independently. There is no interaction between different ions. Therefore

23. For example At 25℃， (NaAc) = 91.0×10-4 S·m2·mol–1，
(HCl)=426.2× S·m2·mol–1，
(NaCl)=126.5× S·m2·mol–1，
What is the molar conductivity of HAc at 25℃?

24. Solution
=( –126.5)×10–4　S·m2·mol–1
=390.7×10–4　 S·m2·mol–1

25. 7.4 Activities of ions and the Debye-Hückel limiting law
Consider an electrolyte
The total chemical potential of an electrolyte is

27. For example

28. mean activity coefficient of ions
Define the mean activity of ions as
mean activity coefficient of ions
mean molality of ions

30. example

32. 7.4.2 The ionic strength The ionic strength, I, is defined by
In dilute solutions, the activity coefficients of electrolytes, the solubilities of sparingly soluble salts, rates of ionic reactions, and other related properties become functions of ionic strength.

33. 3.德拜-休克尔极限公式 (1)离子氛（ionic atmosphere）
这是德拜-休克尔理论中的一个重要概念。他们认为在溶液中，每一个离子都被反号离子所包围，由于正、负离子相互作用，使离子的分布不均匀。
若中心离子取正离子，周围有较多的负离子，部分电荷相互抵消，但余下的电荷在距中心离子 处形成一个球形的负离子氛；反之亦然。一个离子既可为中心离子，又是另一离子氛中的一员。

34. 7.4.3 The Debye-Hückel limiting law

35. In aqueous solution at 25℃
It is only applies to dilute solutions.

37. 7.5 Electrochemical cells
There are two basic types of electrochemical cell.
A galvanic cell and an electrolytic cell

38. This electrochemical cell is represented by the schematic diagram or cell symbol
the vertical lines indicate boundaries between different phases. For boundary between two liquid phases a dashed line ┋ is used to denote. A double dashed line ┋ ┋ used to stand for a salt bridge.
The standard convention is to put the negative pole of the battery on the left hand side and the positive pole on the right hand side.

39. example
Zn|ZnSO4(b)┊CuSO4(b)|Cu

40. example
Pt | H2 (p) | HCl(b) | AgCl(s) |Ag

41. 7.6 Types of electrodes 7.6.1 Metal electrodes
Some time an amalgam instead of a pure metal is used to form a metal electrode.
7.6.2 Gas electrodes
In an alkaline solution,

42. Schematic diagram of hydrogen electrode

43. The oxygen electrode has the same structure as hydrogen electrode
The oxygen electrode has the same structure as hydrogen electrode. The electrode reaction is
In an alkaline solution, it is

44. 7.6.3 Oxidation-reduction electrodes
For example
It is denoted as
Quinhydrone electrode is one of several oxidation-reduction electrodes

45. 7.6.4 Metal-insoluble-salt electrodes
Metal-insoluble-salt electrodes are denoted as M | MX | X-. For example,

46. calomel electrode

47. 7.6.5 Metal-insoluble-oxide electrodes
Metal-insoluble-oxide electrodes consist of a metal covered with one of its insoluble oxide which is in contact with an electrolyte solution.

48. 7.6.6 Ion selective electrodes (ISEs)
The important ion selective electrode is Glass membrane electrode. glass acts as a weak acid
The hydrogen ion activity of the internal solution is held constant. When a solution of different pH from the inside comes in contact with the outside of the glass membrane, the glass is either deprotonated or protonated relative to the inside of the glass.

49. 7.7 Designs of galvanic cells
Example Design a galvanic cell by using following two reactions:
Solution

50. Solution
If oxygen electrode is considered, the reactions for two electrodes are

51. Example Devise galvanic cells with the following diffusion processes.
Solution
(1) Reactions at two electrodes are

52. (2) Solution

53. example Is the following design correct？
This design is not correct, because two ions, hydrogen ion and hydroxyl ion, are in different solutions.

54. The correct design is

55. 7.8 Electrode potentials 7.8.1 The EMF in a galvanic cell
If the cell is connected to a measurement circuit, an inevitable potential difference called contact potential between two electrodes, copper and zinc in the Daniell cell, exists as a part of total EMF of the cell.
Cu|Zn|ZnSO4(a1)|CuSO4(a2)|Cu

56. is the potential between solution of copper sulfate and electrode of copper;
is the potential between solution of zinc sulfate and electrode of zinc;
is the potential between two solutions of copper sulfate and zinc sulfate, it is called the liquid junction potential;
is the contact potential between two metals of zinc and copper.

57. 7.8.2 Potential difference between electrode and solution

58. 7.8.3 Potential difference between two solutions
For a simple case of I-I type of electrolyte,

59. 7.8.4 Potential of standard hydrogen electrode
The ultimate reference is the normal hydrogen electrode (NHE), or the standard hydrogen electrode (SHE), the potential of which is defined as zero at all temperatures.
The standard hydrogen electrode consists of a platinum electrode in contact with a solution of H+ at unit activity and saturated with H2 gas with the standard pressure of 100 kPa.

60. 7.8.5 Standard electrode potentials

61. 7.9 Electrochemical thermodynamics
7.9.1 Gibbs function and reversible EMF
EMF: Electromotive force
For a cell reaction at constant temperature and pressure,

62. 7.9.2 Entropy and enthalpy of cell reaction
at constant temperature,
In a reversible galvanic cell

63. For a reversible galvanic cell at constant pressure
The reversible heat in a galvanic cell is different from the heat of chemical reaction

64. Example
The EMF and temperature coefficient are V and -5.00×10-5V·K-1 respectively for the Weston standard cell at 25℃. Calculate the
and Qr.
Solution The cell reaction is
Anode
Cathode
Cell

66. Weston standard cell

67. 7.9.3 The Nernst equation
For a chemical reaction or an electrochemical cell reaction, the isotherm equation is
25℃, and convert natural logarithm into common logarithm, equation can be expressed as

68. When cell reaction reaches the equilibrium state, E=0, then
The Nernst equation applies not only to the electrochemical cells but also to the electrodes.
For instance, Cl2(g) + 2e = 2Cl-

69. Cu2++2e-→Cu
example
example O2(g) + 4H+ + 4e–  2H2O

70. Calculation of EMF
Total:
method 1:

71. Method 2:
total:
Two methods, one result.

72. example
Write the reactions of the following cell and calculate the EMF at 25℃ when b(HCl)=0.1mol kg-1.
Solution
Anode:
Cathode:
Total:

73. From table 7.4.1, the activity coefficient can be obtained. therefore
From this equation we can conclude that activity coefficient can be obtained in terms of the determination of EMF of the cell at different concentrations of HCl.

74. 7.9.4 Determination of E values
Take the following cell as an example

75. For small enough values of b, γ± will be given by the Debye-Hückel formula. For HC1, the ionic strength is equal to the molality b if no other electrolytes are present. Therefore,
the left-hand side of this equation is plotted as a function of b1/2, the extrapolation to b = 0 gives the value of E from the intercept.

77. 7.9.5 Pourbaix diagrams Marcel Pourbaix (1904-1998)
Marcel Pourbaix was born in Russia, where his father, a Belgian engineer, was working at the time.
The significance of Marcel Pourbaix’s great achievement　was pointed out by Ulick R. Evans, widely recognized as the　“father of corrosion science,” in his foreword to Pourbaix’s　Thermodynamics of Dilute Aqueous Solutions: “During the last　decade (the 1940s) Dr. Marcel Pourbaix of Brussels has developed a graphical method, based on generalized thermodynamical equations, for the solution of many different kinds of scientific problems, involving numerous types of heterogeneous or homogeneous reactions and equilibria... Some of these problems have long been treated from the aspect of thermodynamics... The application of thermodynamics to typical corrosion reactions is a much newer development.”

78. For Mr. Zhang Zhongcheng
As a memory of his stay at Cebelcor in 1984/1985.
With all best wishes from
　　　　　　　　Marcel Pourbaix
　　　　　　　　　April 26, 1985

79. Prof Pourbaix, his wife, and Zhang Zhongcheng, in 1985

80. Prof Pourbaix with his colleagues at Cebelcor in 1984

81. The E-pH diagram of copper-water system

82. 7.10 Decomposition voltage

83. 7.10 Decomposition voltage氯气 氢气

84. Decomposition voltage is a cell voltage at which the electrolysis current begins to increase appreciably.
evolution potentials

85. 7.11 Polarization 7.11.1 Electrode polarization
The deviation of voltage from the value at equilibrium upon the passage of current is called polarization.
For an electrode, the deviation of electrode potential from its equilibrium state is called the electrode polarization.
overpotential, denoted as η and expressed in volts.

86. Various overpotentials
Concentration overpotential
Activation overpotential (electrochemical overpotential)
Ohmic polarization

87. 7.11.2 Determination of polarization curves

88. Polarization curves
(Tafel equation)

89. 7.11.3 Polarizations of electrolytic and galvanic cells

90. 7.12 Reactions in electrolytic cells
At anode, the reaction which has the lowest polarized potential E(anode) takes place.
At cathode, the reaction which has the highest polarized potential E(cathode) takes place.

91. Example
A zinc cathode is used to electrolyze an aqueous solution of ZnSO4 (a±=1). What will give off at cathode under atmospheric pressure, hydrogen or zinc? The overpotential of hydrogen on zinc is 0.7V.
Solution The overpotential of zinc on the cathode is negligible.

92. The evolution potential of hydrogen at equilibrium on cathode is
Suppose the pressure of hydrogen is kPa, the aqueous solution is neutral, a(H+)=10-7. Then,
Consider the overpotential of hydrogen on zinc, the polarized potential of hydrogen evolution is
This value is lower than the deposition potential of zinc. Therefore, it is zinc that forms on the cathode.

93. 7.13 Corrosion
Ulick R. Evans, the British scientist who is considered the "Father of Corrosion Science", has given the definition of corrosion:
"Corrosion is largely an electrochemical phenomenon, may be defined as destruction by electrochemical or chemical agencies".

94. Fig 7.13.2 Schematic representation of current flow in a simple corrosion cell

95. Fig 7.13.3 The galvanic corrosion of zinc plate caused by Fe impurity

96. Corrosion
19.7

97. Dental Filling Discomfort
Hg2 /Ag2Hg V 2+
Sn /Ag3Sn V 2+

98. Fig 7.13.1 Simplified E-pH diagram of Fe-H2O system

99. Measures preventing metals from corrosion
Cathodic protection
There are two ways to carry out a cathode protection.
One is the attachment of an active metal to a system to protect the system from corrosion.
For example, galvanized iron.
Another way to carry out cathode protection is the application of an external voltage or current to set the voltage of the material at a sufficiently negative potential such that the rate of corrosion is very low.

100. Cathodic Protection of an Iron Storage Tank
19.7

101. Anode protection

102. Batteries Dry cell Leclanché cell Zn (s) Zn2+ (aq) + 2e- Anode:
Cathode:
2NH4 (aq) + 2MnO2 (s) + 2e Mn2O3 (s) + 2NH3 (aq) + H2O (l) +
Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)

104. Batteries Mercury Battery Anode:
Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-
Cathode:
HgO (s) + H2O (l) + 2e Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)

105. Batteries Lead storage battery Anode:
Pb (s) + SO2- (aq) PbSO4 (s) + 2e- 4
Cathode:
PbO2 (s) + 4H+ (aq) + SO2- (aq) + 2e PbSO4 (s) + 2H2O (l) 4
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) PbSO4 (s) + 2H2O (l) 4

106. Batteries
A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning
Anode:
2H2 (g) + 4OH- (aq) H2O (l) + 4e-
Cathode:
O2 (g) + 2H2O (l) + 4e OH- (aq)
2H2 (g) + O2 (g) H2O (l)