Electrochemistry

1. Electrochemistry Dr. ShuchitaAgrawal BTIRT Sironja, Sagar

2. Electrochemistry Electrolysis Electric energyChemical energy Galvanic cell

3. Definition of electrochemistry I • A science studying the relationship between chemical energyand 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.

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

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

6. Faraday's lawof 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 = 96485.309 C·mol-1= 1 Faraday.

7. Coulometer • The silver coulometer and copper coulometer are commonly used ones. • If one mole of silver, 107.870 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.

8. Transport numbers and mobilities • The definition of transport numbert 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 numberof 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.

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

10. 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. Conductance, conductivity and molar conductivity • 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 • ExampleThe conductivity and molar conductivity of a saturated aqueous solution of silver chloride are 3.41×10-4S·m-1 and 138.26×10-4S·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. Solutionthe 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. The dependence of molar conductivity on the concentration For strong electrolytes, Kohlrauschobserved that m decreased with concentration according to the expression

20. conductivities and the concentrations of electrolytes

21. 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)=426.2×10-4 S·m2·mol–1， (NaCl)=126.5×10-4 S·m2·mol–1， What is the molar conductivity of HAc at 25℃?

24. Solution =(426.3 +91.0–126.5)×10–4　S·m2·mol–1 =390.7×10–4　 S·m2·mol–1

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

27. For example

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

29. Example

31. 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.

32. The Debye-Hückel limiting law

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

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

36. 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.

37. Example Zn|ZnSO4(b)┊CuSO4(b)|Cu

38. Example Pt | H2 (p) | HCl(b) | AgCl(s) |Ag

39. Types of electrodes • Metal electrodes Zn2+|Zn: Zn2+ + 2e → Zn • Some time an amalgam instead of a pure metal is used to form a metal electrode. • Gas electrodes • In an alkaline solution,2H+ + 2e → H2 2H2O + 2e → H2(g) + 2OH-

40. Hydrogen electrode

41. The oxygen electrode has the same structure as hydrogen electrode. The electrode reaction is O2(g) + 4H+ + 4e → 2H2O • In an alkaline solution, it is O2(g) + 2H2O + 4e → 4OH-

42. Oxidation-reduction electrodes • For example • Fe3+ + e→ Fe2+ It is denoted as • Quinhydrone electrode is one of several oxidation-reduction electrodes C6H4O2 + 2H+ + 2e → C6H4(OH)2

43. Metal-insoluble-salt electrodes • Metal-insoluble-salt electrodes are denoted as M | MX | X-. For example, AgCl (s) + e →Ag(s) + Cl-

44. Calomel electrode Hg2Cl2 (s) + 2e → 2Hg + 2Cl-

45. 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. , Sb2O3(s) + 3H2O+ 6e → 2Sb + 6OH-

46. Ion selective electrodes (ISEs) • The important ion selective electrode is Glass membrane electrode. glass acts as a weak acid Glass – H → Glass + H+ 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 de-protonated or protonated relative to the inside of the glass.

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

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

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

50. (2) Solution