HCM INTERNATIONAL UNIVERSITY SCHOOL OF BIOTECHNOLOGY

1. HCM INTERNATIONAL UNIVERSITYSCHOOL OF BIOTECHNOLOGY APPLICATIONS OF ELECTROCHEMISTRY Course: ANALYTICAL CHEMISTRY Lecturer: Dr. NGUYEN TUAN KHOI

2. MEMBERS OF GROUP Ph ạm Nguyễn Huệ Nh ân Thái Văn Ch í Nguyễn Th ị Phương Thùy Trần Đỗ Ngọc Oanh V õ Hoàng Lâm Nguyễn Th ị Thu Cúc Trương Thị Ngọc Nhi Lê Trần Khánh Trang Đỗ Vân Khanh T ôn Th ị H ồng Th ảo Nguy ễn Vi ệt Th ư Nguy ễn Ng ọc Y ên Nhi Đ oàn T ây Nguy ên Nguy ễn Duy Trung Nguy ễn Đ ức Thanh Long V ũ Ng ọc C ư ơng Nguyễn Vũ Nh ất Th ịnh

3. Outline I. Introduction II. Electrochemical cells • Galvanic cells • Electrolytic cells III. Electrochemical cell applications • Battery • Corrosion • Electrolysis IV. Electrochemical methods • Nernst equation • Potentiometry • Coulometry • Voltammetry

4. I. INTRODUCTION • Electrochemistry is the study of reactions in which charged particles (ions or electrons) cross the interface between two phases of matter, typically a metallic phase (the electrode) and a conductive solution, or electrolyte. This reaction is simple oxidation-reduction process.

5. I. INTRODUCTION Redox reaction (reduction-oxidation reactions) • Are reactions that mention to the transfer of electrons between species. • Describe all chemical reactions in which atoms have oxidation number change. Ox1 + red2  red1 + ox2

6. II. Electrochemical cells

8. II. Electrochemical cell • Transform energy from chemical reaction to electrical energy or vice versa. • An electrochemical cell includes: • Two electrodes: half redox reactions occur • Anode: oxidation reaction occur • Cathode: reduction reaction occur • Electrolyte solution(s)

9. II. Electrochemical cell • Conditions for generating electricity flow: • The electrodes must be externally connected by a metal wire to permit electron flow. • The electrolyte solutions are in contact to allow movement of ions.

10. II. Electrochemical cell • There are two types of electrochemical cells: • Galvanic cells (or Voltaic cells): spontaneous reactions occur. • Electrolytic cells: nonspontaneous reactions occur (electrical energy supply).

11. a. GALVANIC CELL What about the sign of the electrodes? What about half-cell reactions? - + Why? cathode half-cell Cu+2 + 2e- Cu anode half-cell Zn  Zn+2 + 2e- Cu plates out or deposits on electrode Zn electrode erodes or dissolves What happened at each electrode? Cu Zn 1.0 M CuSO4 1.0 M ZnSO4

12. Zn (s) + Cu2+(aq) Cu (s) + Zn2+(aq) a. Galvanic cells HALF REACTION Anode (Ox) : Zn(s) = Zn2+ + 2e Cathode (Red) : Cu2+ + 2e = Cu (s) Net reaction : Zn (s) + Cu2+ = Zn2+ +Cu (s) Salt bridge Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s) cathode anode

13. b. ELECTROLYTIC CELL

14. CELL POTENTIAL • cell potential: Electrons flow from one electrode to the other in one direction, there is a potential difference between the electrodes. • Cell potential is calculated in voltage (V) by the formula: • E cathode: reduction potential (V) • E anode: oxidation potential (V)

15. III. Electrochemical cell applications

18. PRIMARY BATTERY Primary battery has long been known as dry-cell. It cannot be recharged. It’s widely used to power flashlight and some other similar devices. The first practical battery consisted of a stack of small electrical cell, each consisting of a silver plate and a zinc plate separated by a sheet of cardboard which had been soaked in salt water

19. A TYPICAL STRUCTURE OF A PRIMARY BATTERY The electrode reactions are Zn → Zn2+ + 2e– 2 MnO2 + 2H+ + 2e– → Mn2O3 + H2O

20. SECONDARY BATTERIES A secondaryorstorage battery is capable of being recharged. Its electrode reactions can proceed in either direction. During charging, electrical work is done on the cell to provide the free energy needed to force the reaction in the non-spontaneous direction.

22. Fuel cell Conventional batteries supply electrical energy from the chemical reactants stored within them. When these reactants are consumed, the battery is "dead". An alternative approach would be to feed the reactants into the cell as they are required, so as to permit the cell to operate continuously. In this case the reactants can be thought of as "fuel" to drive the cell, hence the term fuel cell.

23. MODERN HYDROGEN-OXYGEN FUEL CELL

25. ELECTROCHEMICAL CORROSION Corrosion is the deterioration of materials by chemical processes. Of these, the most important by far is electrochemical corrosion of metals, in which the oxidation process M → M+ + e– is facilitated by the presence of a suitable electron acceptor

28. Sacrificial coating One way of supplying this negative charge is to apply a coating of a more active metal a very common way of protecting steel from corrosion is to coat it with a thin layer of zinc

30. Cathodic protection A more sophisticated strategy is to maintain a continual negative electrical charge on a metal, so that its dissolution as positive ions is inhibited. The entire surface is forced into the cathodic condition.

31. ELECTROLYSIS Electrolysis refers to the decomposition of a substance by an electric current

32. ELECTROLYSIS OF WATER • Water is capable of undergoing both oxidation and reduction • Pure water is an insulator and cannot undergo signifigant electrolysis without adding an electrolyte. • Electrolysis of a solution of sulfuric acid or of a salt such as NaNO3 results in the decomposition of water at both electrodes: • cathode:  H2O + 2 e– → H2(g) + 2 OH– E =+0.41 v ([OH–] = 10-7 M) • anode:  2 H2O → O2(g) + 4 H+ + 2 e– E° = -0.82 v • net:   2 H2O(l) → 2 H2(g) + O2(g) E = -1.23 v

33. ELECTROLYSIS

34. THE CHLORALKALI INDUSTRY • The electrolysis of brine is carried out on a huge scale for the industrial production of chlorine and caustic soda (sodium hydroxide). Because the reduction potential of Na+ is much higher than that of water, the latter substance undergoes decomposition at the cathode, yielding hydrogen gas and OH–. • 2 NaCl + 2 H2O → 2 NaOH + Cl2(g) + H2(g)

35. A modern industrial chloralkali plant

36. ELECTROLYTIC REFINING OF ALUMINUM • The Hall-Hérault process takes advantage of the principle that the melting point of a substance is reduced by admixture with another substance with which it forms a homogeneous phase. • The net reaction is 2 Al2O3 + 3 C → 4 Al + 3 CO2

37. IV. Electroanalytical methods

38. NERNST EQUATION Nernst equation allows one unknown concentration to be determined from a measurement of the cell voltage. aOx + ne- ↔ bRed E = E0 – (2.3026RT/nF)log ([Red]b/[Ox]a) E: the reduction potential at the specific concentration n: the number of electrons R: the gas constant (8.3143 V coul deg-1 mol -1) T: the absolute temperature F: the Faraday constant (96487 colul eq-1) At 25oC, the value of 2.3026RT/F is 0.05916

39. ELECTROCHEMICAL ANALYSIS • Potentiometry • Coulometry • Voltammetry

40. a. POTENTIOMETRY • Potentiometry passively measures the potential of a solution between two electrodes, affecting the solution very little in the process. The potential is then related to the concentration of one or more analytes. • In potentiometry, there are no current, or only negligible current flows, so the compound in the solution remain unchanged. It is used for measure the cell potential and for determine the analytical quantity of interest. Potentiometry is a useful quantitative method.

41. Potentiometric titration

42. Mechanism Ecell= Eind - Eref Difference in potential Reference electrode Indicator electrode (Constant potential) (Change in potential) Mobilities of ions Solution

43. Electrode Metal Indicator Electrode Membrane Potentiometry Electrode Reference Electrode

44. Reference electrodes Calomel Reference Electrodes Silver/ Silver Chloride Reference Electrodes

45. Indicator electrodes Metallic Membrane

46. Application Used in pH meter, by using glass electrode In environment, used to analyse ion -CN-, F-, NH3, and NO3- in water and in wastewater.

47. b. COULOMETRY Coulometry: electrochemical method based on the quantitative oxidation or reduction of analyte - Measure amount of analyte by measuring amount of current and time required to complete reaction charge = current (i) x time in coulombs Q = ite

48. Application of coulometry

49. A coulometer is a device used for measuring the quantity of electricity required to bring about a chemical change of the analyte. It is usual practice in coulometry to substitute the ammeter COULOMETER:

50. c. Voltammetry • Measures current as a function of applied potential under conditions that keep a working electrode polarized