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Presentation On Electrochemistry

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Published in: Chemistry
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Electro Chemistry ppt part I

Sunil / Bangalore

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Qualification: B.Tech/B.E. (Kolhapur Institute of Technology's College of Engineering (KITCE), Kolhapur - 2018)

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  1. Electrochemistry Part - I
  2. Electrochemistry Electrochemistry is the study of production of electricity from the energy released during a spontaneous chemical reaction and the use of electrical energy to bring about non-spontaneous chemical transformations. Electrochemical cells A spontaneous chemical process is the one which can take place on its own and in such a process the Gibb's energy of the system decreases. It is this energy that gets converted to electrical energy. The reverse process is also possible in which we can make non-spontaneous processes occur by supplying external energy in the form of electrical energy. These inter conversions are carried out in equipment's called Electrochemical Cells.
  3. Types Electrochemical Cells are of two types: 1. Galvanic Cells Converts chemical energy into electrical energy 2. Electrolytic Cells Converts electrical energy into chemical energy.
  4. Electrochemistry Part - I
  5. 1. Galvanic cell Cell energy is extracted from a spontaneous chemical process or reaction and it is converted to electric current. For example, Daniel Cell is a Galvanic Cell in which Zinc and Copper are used for the redox reaction to take place. cu2+ -Y Zn2+ + Cucs) (aq) (aq) Oxidation Half : Zn (s) -->Zn2+ (aq) Reduction Half : Cu2+ + 2e- Cu (aq) (s) Zn is the reducing agent and Cu2+ is the oxidizing agent. The half cells are also known as Electrodes. The oxidation half is known as Anode and the reduction half is called Cathode. Electrons flow from anode to cathode in the external circuit. Anode is assigned negative polarity and cathode is assigned positive polarity. In Daniel Cell, Zn acts as the anode and Cu acts as the cathode.
  6. Electrochemistry Electrochemistry is the study of production of electricity from the energy released during a spontaneous chemical reaction and the use of electrical energy to bring about non-spontaneous chemical transformations. Electrochemical cells A spontaneous chemical process is the one which can take place on its own and in such a process the Gibb's energy of the system decreases. It is this energy that gets converted to electrical energy. The reverse process is also possible in which we can make non-spontaneous processes occur by supplying external energy in the form of electrical energy. These inter conversions are carried out in equipment's called Electrochemical Cells.
  7. Electrolytic cell These electrodes are dipped in and electrolytic solution containing cations and anions. On supplying current the ions move towards electrodes of opposite polarity and simultaneous reduction and oxidation takes place. Preferential Discharge of ions Where there are more than one cation or anion the process of discharge becomes competitive in nature. Discharge of any ion requires energy and in case of several ions being present the discharge of that ion will take place first which requires the energy.
  8. Types Electrochemical Cells are of two types: 1. Galvanic Cells Converts chemical energy into electrical energy 2. Electrolytic Cells Converts electrical energy into chemical energy.
  9. Electrode potential It may be defined as the tendency of an element, when it is placed in contact with its own ions to either lose or gain electrons and in turn become positively or negatively charged. The electrode potential will be named as oxidation or reduction potential depending upon whether oxidation or reduction has taken place. (aq) Reduction Oxidation Characteristics (a) Both oxidation and reduction potentials are equal in magnitude but opposite in sign. (b) It is not a thermodynamic property, so values of E are not additive.
  10. 1. Galvanic cell Cell energy is extracted from a spontaneous chemical process or reaction and it is converted to electric current. For example, Daniel Cell is a Galvanic Cell in which Zinc and Copper are used for the redox reaction to take place. cu2+ -Y Zn2+ + Cucs) (aq) (aq) Oxidation Half : Zn (s) -->Zn2+ (aq) Reduction Half : Cu2+ + 2e- Cu (aq) (s) Zn is the reducing agent and Cu2+ is the oxidizing agent. The half cells are also known as Electrodes. The oxidation half is known as Anode and the reduction half is called Cathode. Electrons flow from anode to cathode in the external circuit. Anode is assigned negative polarity and cathode is assigned positive polarity. In Daniel Cell, Zn acts as the anode and Cu acts as the cathode.
  11. Standard electrode potential (EO) It may be defined as the electrode potential of an electrode determined relative to standard hydrogen electrode under standard conditions. The standard conditions taken are 1. 1M concentration of each ion in the solution. 2. A temperature of 298 K. 3. 1 bar pressure for each gas.
  12. Electrolytic cell These electrodes are dipped in and electrolytic solution containing cations and anions. On supplying current the ions move towards electrodes of opposite polarity and simultaneous reduction and oxidation takes place. Preferential Discharge of ions Where there are more than one cation or anion the process of discharge becomes competitive in nature. Discharge of any ion requires energy and in case of several ions being present the discharge of that ion will take place first which requires the energy.
  13. Electrochemical series The half cell potential values are standard values and are represented as the standard reduction potential values as shown in the table at the end which is also called Electrochemical Series. Stronger oxidizing agent Weaker oxidizing agent F2(S) + 2 e— H202(aq) + 2 H+(aq) + 2e- C12Cg) + 2 e- + 14 H+(aq) + 6 02(g) + 4 H+(aq) + 4 e Br2(/) + 2 e- Ag+(aq) + e— Fe3+(aq) + e— 02(S) + 2 (aq) + 2 e 12(s) + 2 e— 02(S) + 2 H20(1) Cu2+(aq) + 2 e— Sn4+(aq) + 2 e- Pb2+(aq) + 2e- Ni2+(aq) + 2e- Zn2+(aq) + 2c- A13+(aq) + 3 e- Mg2+(aq) + 2 e— Li+(aq) + e— -...........-.> 2 F (aq) .-.......-.-..> 2 cr3+(aq) + 7 H20(/) .....................-> Ag(S) -...........-> Fe2+(aq) —.............-> H202(aq) ...........-...> Cu(s) ......-........> sn2+(aq) -........-..> H2(g) -.....-...-...> Pb(s) ...............-..> Ni(s) .................-> Fe(s) -......--> Zn(s) ..................> 112Cg) ..........-.-.> Al(s) -........-.-..> Mg(s) -......-......-.> Na(s) + 2 OH-(aq) 2.87 1.78 1.51 1.36 133 1.23 1.09 0.80 0.77 0.70 0.54 0.40 034 0.15 -0.13 -026 -0.40 -0.45 -0.76 -0.83 —1.66 -2.37 -2.71 -3.04 Wea ker reducing agent Stronger red ucing agent
  14. Electrode potential It may be defined as the tendency of an element, when it is placed in contact with its own ions to either lose or gain electrons and in turn become positively or negatively charged. The electrode potential will be named as oxidation or reduction potential depending upon whether oxidation or reduction has taken place. (aq) Reduction Oxidation Characteristics (a) Both oxidation and reduction potentials are equal in magnitude but opposite in sign. (b) It is not a thermodynamic property, so values of E are not additive.
  15. Cell potential or emf of a cell The difference between the electrode potentials of two half cells is called cell potential. It is known as electromotive force (EMF) of the cell if no current is drawn from the cell. cell cathode anode For this equation we take oxidation potential of anode and reduction potential of cathode. Since anode is put on left and cathode on right, it follows therefore, Left Right cell For a Daniel cell, therefore 2 = 0.34+0.76 = 1.10 v cell — Zn + cu
  16. Standard electrode potential (EO) It may be defined as the electrode potential of an electrode determined relative to standard hydrogen electrode under standard conditions. The standard conditions taken are 1. 1M concentration of each ion in the solution. 2. A temperature of 298 K. 3. 1 bar pressure for each gas.
  17. Cell diagram or representation of a cell The following conventions or notations are applied for writing the cell diagram in accordance with IUPAC recommendations. The Daniel cell is represented as follows: zn(s) I Zn2+(C1) Il cu2+(C2) I cucs) Anode half cell is written on the left hand side while cathode half cell on 1. 1. 2. 3. right hand side. zn(s) I Zn2+ (aq) ; Anodic chamber cu2+ I Cucs) (aq) Cathodic chamber A single vertical line separates the metal from aqueous solution of its own ions. A double vertical line represents salt bridge The molar concentration (C) is placed in brackets after the formula of the corresponding ion.
  18. Electrochemical series The half cell potential values are standard values and are represented as the standard reduction potential values as shown in the table at the end which is also called Electrochemical Series. Stronger oxidizing agent Weaker oxidizing agent F2(S) + 2 e— H202(aq) + 2 H+(aq) + 2e- C12Cg) + 2 e- + 14 H+(aq) + 6 02(g) + 4 H+(aq) + 4 e Br2(/) + 2 e- Ag+(aq) + e— Fe3+(aq) + e— 02(S) + 2 (aq) + 2 e 12(s) + 2 e— 02(S) + 2 H20(1) Cu2+(aq) + 2 e— Sn4+(aq) + 2 e- Pb2+(aq) + 2e- Ni2+(aq) + 2e- Zn2+(aq) + 2c- A13+(aq) + 3 e- Mg2+(aq) + 2 e— Li+(aq) + e— -...........-.> 2 F (aq) .-.......-.-..> 2 cr3+(aq) + 7 H20(/) .....................-> Ag(S) -...........-> Fe2+(aq) —.............-> H202(aq) ...........-...> Cu(s) ......-........> sn2+(aq) -........-..> H2(g) -.....-...-...> Pb(s) ...............-..> Ni(s) .................-> Fe(s) -......--> Zn(s) ..................> 112Cg) ..........-.-.> Al(s) -........-.-..> Mg(s) -......-......-.> Na(s) + 2 OH-(aq) 2.87 1.78 1.51 1.36 133 1.23 1.09 0.80 0.77 0.70 0.54 0.40 034 0.15 -0.13 -026 -0.40 -0.45 -0.76 -0.83 —1.66 -2.37 -2.71 -3.04 Wea ker reducing agent Stronger red ucing agent
  19. 4. 5. The value of e.m.f. of the cell is written on the extreme right of the cell. For example, zn(s) I Zn2+(1M) Il Cu2+(1M) I cucs) E.M.F = +1.1 V If an inert electrode like platinum is involved in the construction of the cell, it may be written along with the working electrode in bracket say for example, when a zinc anode Is connected to a hydrogen electrode. zn(s) I Zn2+(C1) Il H+(C2) I H2
  20. Cell potential or emf of a cell The difference between the electrode potentials of two half cells is called cell potential. It is known as electromotive force (EMF) of the cell if no current is drawn from the cell. cell cathode anode For this equation we take oxidation potential of anode and reduction potential of cathode. Since anode is put on left and cathode on right, it follows therefore, Left Right cell For a Daniel cell, therefore 2 = 0.34+0.76 = 1.10 v cell — Zn + cu
  21. Salt bridge 1. 2. 3. Salt bridge is used to maintain the charge balance and to complete the circuit by facilitating the flow of ions through it. It contains a gel in which an inert electrolyte like Na2S04 or KN03 etc are mixed. Negative ions flow to the anode and positive ions flow to the cathode through the salt bridge and charge balance is maintained and cell keeps on functioning.
  22. Cell diagram or representation of a cell The following conventions or notations are applied for writing the cell diagram in accordance with IUPAC recommendations. The Daniel cell is represented as follows: zn(s) I Zn2+(C1) Il cu2+(C2) I cucs) Anode half cell is written on the left hand side while cathode half cell on 1. 1. 2. 3. right hand side. zn(s) I Zn2+ (aq) ; Anodic chamber cu2+ I Cucs) (aq) Cathodic chamber A single vertical line separates the metal from aqueous solution of its own ions. A double vertical line represents salt bridge The molar concentration (C) is placed in brackets after the formula of the corresponding ion.
  23. Zinc Sol, VolimctctJAmme1cr soi 'bridge —???• soiz—-——-—- —???????— CillElUdi2 Coppcr Sol. Anode ???????#?• Cu —+ CII zn-cuso,-.zr.scu-cu
  24. 4. 5. The value of e.m.f. of the cell is written on the extreme right of the cell. For example, zn(s) I Zn2+(1M) Il Cu2+(1M) I cucs) E.M.F = +1.1 V If an inert electrode like platinum is involved in the construction of the cell, it may be written along with the working electrode in bracket say for example, when a zinc anode Is connected to a hydrogen electrode. zn(s) I Zn2+(C1) Il H+(C2) I H2
  25. Spontaneity of a reaction AG = - nFE Cell 1. For a spontaneous cell reaction AG should be negative and cell potential should be positive. 2. If we take standard value of cell potential in the above equation we will obtain standard value of AG as well. i.e. AGO - nFEO Cell
  26. Salt bridge 1. 2. 3. Salt bridge is used to maintain the charge balance and to complete the circuit by facilitating the flow of ions through it. It contains a gel in which an inert electrolyte like Na2S04 or KN03 etc are mixed. Negative ions flow to the anode and positive ions flow to the cathode through the salt bridge and charge balance is maintained and cell keeps on functioning.
  27. Types of electrodes 1. Metal-Metal Ion electrodes A metal rod/plate is dipped in an electrolyte solution containing metal ions. There is a potential difference between these two phases and this electrode can act as a cathode or anode both. Anode: M Mn++ ne- Cathode: Mll+ + ne- M
  28. Zinc Sol, VolimctctJAmme1cr soi 'bridge —???• soiz—-——-—- —???????— CillElUdi2 Coppcr Sol. Anode ???????#?• Cu —+ CII zn-cuso,-.zr.scu-cu
  29. 2. Gas Electrodes Electrode gases like H2, C12 etc are used with their respective ions. For example, H2 gas is used with a dilute solution of HCI (H + ions). The metal should be inert so that it does not react with the acid. Anode: 2H++ ze- Cathode: 2H++ ze- The hydrogen electrode is also used as the standard to measure other electrode potentials. Its own potential is set to 0 V as a reference. When it is used as a reference the concentration of dil. HCI is taken as 1 M and the electrode is called "Standard Hydrogen Electrode (SHE)".
  30. Spontaneity of a reaction AG = - nFE Cell 1. For a spontaneous cell reaction AG should be negative and cell potential should be positive. 2. If we take standard value of cell potential in the above equation we will obtain standard value of AG as well. i.e. AGO - nFEO Cell
  31. Types of electrodes 1. Metal-Metal Ion electrodes A metal rod/plate is dipped in an electrolyte solution containing metal ions. There is a potential difference between these two phases and this electrode can act as a cathode or anode both. Anode: M Mn++ ne- Cathode: Mll+ + ne- M
  32. ? ? ? ? ? ? ? t ? ? ? ? ? ? 1 0 0 ? With ?
  33. 2. Gas Electrodes Electrode gases like H2, C12 etc are used with their respective ions. For example, H2 gas is used with a dilute solution of HCI (H + ions). The metal should be inert so that it does not react with the acid. Anode: 2H++ ze- Cathode: 2H++ ze- The hydrogen electrode is also used as the standard to measure other electrode potentials. Its own potential is set to 0 V as a reference. When it is used as a reference the concentration of dil. HCI is taken as 1 M and the electrode is called "Standard Hydrogen Electrode (SHE)".
  34. 3. Metal-Insoluble salt electrode We use salts of some metals which are sparingly soluble with the metal itself as electrodes. For example, if we use AgCl with Ag there is a potential gap between these two phases which can be identified in the following reaction: Agcl(s) + e- + cl- This electrode is made by dipping a silver rod in a solution containing AgCl(s) and Cl- ions.
  35. ? ? ? ? ? ? ? t ? ? ? ? ? ? 1 0 0 ? With ?
  36. 4. Calomel Electrode Mercury is used with two other phases, one is a calomel paste (Hg2C12) and electrolyte containing Cl- ions. Cathode : ze- 2Hg(l) + zcl- (aq) Anode : This electrode is also used as another standard to measure other potentials. Its standard form is also called Standard Calomel Electrode (SCE).
  37. Electrochemistry Part - I
  38. 3. Metal-Insoluble salt electrode We use salts of some metals which are sparingly soluble with the metal itself as electrodes. For example, if we use AgCl with Ag there is a potential gap between these two phases which can be identified in the following reaction: Agcl(s) + e- + cl- This electrode is made by dipping a silver rod in a solution containing AgCl(s) and Cl- ions.
  39. TO salt bridge pt. Wire Saturated KCt solution EEh'2Cin and Hg paste Bletx:uy Calomel electrode
  40. Electrochemistry Electrochemistry is the study of production of electricity from the energy released during a spontaneous chemical reaction and the use of electrical energy to bring about non-spontaneous chemical transformations. Electrochemical cells A spontaneous chemical process is the one which can take place on its own and in such a process the Gibb's energy of the system decreases. It is this energy that gets converted to electrical energy. The reverse process is also possible in which we can make non-spontaneous processes occur by supplying external energy in the form of electrical energy. These inter conversions are carried out in equipment's called Electrochemical Cells.
  41. 4. Calomel Electrode Mercury is used with two other phases, one is a calomel paste (Hg2C12) and electrolyte containing Cl- ions. Cathode : ze- 2Hg(l) + zcl- (aq) Anode : This electrode is also used as another standard to measure other potentials. Its standard form is also called Standard Calomel Electrode (SCE).
  42. 5. Redox Electrode In these electrodes two different oxidation states of the same metal are used in the same half cell. For example, Fe2+ and Fe3+ are dissolved in the same container and an inert electrode of platinum is used for the electron transfer. Following reactions can take place: Anode: Fe2+ Fe3+ + e- Cathode: Fe3+ + e- Fe2+
  43. Types Electrochemical Cells are of two types: 1. Galvanic Cells Converts chemical energy into electrical energy 2. Electrolytic Cells Converts electrical energy into chemical energy.
  44. TO salt bridge pt. Wire Saturated KCt solution EEh'2Cin and Hg paste Bletx:uy Calomel electrode
  45. 1. Galvanic cell Cell energy is extracted from a spontaneous chemical process or reaction and it is converted to electric current. For example, Daniel Cell is a Galvanic Cell in which Zinc and Copper are used for the redox reaction to take place. cu2+ -Y Zn2+ + Cucs) (aq) (aq) Oxidation Half : Zn (s) -->Zn2+ (aq) Reduction Half : Cu2+ + 2e- Cu (aq) (s) Zn is the reducing agent and Cu2+ is the oxidizing agent. The half cells are also known as Electrodes. The oxidation half is known as Anode and the reduction half is called Cathode. Electrons flow from anode to cathode in the external circuit. Anode is assigned negative polarity and cathode is assigned positive polarity. In Daniel Cell, Zn acts as the anode and Cu acts as the cathode.
  46. 5. Redox Electrode In these electrodes two different oxidation states of the same metal are used in the same half cell. For example, Fe2+ and Fe3+ are dissolved in the same container and an inert electrode of platinum is used for the electron transfer. Following reactions can take place: Anode: Fe2+ Fe3+ + e- Cathode: Fe3+ + e- Fe2+
  47. Electrolytic cell These electrodes are dipped in and electrolytic solution containing cations and anions. On supplying current the ions move towards electrodes of opposite polarity and simultaneous reduction and oxidation takes place. Preferential Discharge of ions Where there are more than one cation or anion the process of discharge becomes competitive in nature. Discharge of any ion requires energy and in case of several ions being present the discharge of that ion will take place first which requires the energy.
  48. Electrode potential It may be defined as the tendency of an element, when it is placed in contact with its own ions to either lose or gain electrons and in turn become positively or negatively charged. The electrode potential will be named as oxidation or reduction potential depending upon whether oxidation or reduction has taken place. (aq) Reduction Oxidation Characteristics (a) Both oxidation and reduction potentials are equal in magnitude but opposite in sign. (b) It is not a thermodynamic property, so values of E are not additive.
  49. Standard electrode potential (EO) It may be defined as the electrode potential of an electrode determined relative to standard hydrogen electrode under standard conditions. The standard conditions taken are 1. 1M concentration of each ion in the solution. 2. A temperature of 298 K. 3. 1 bar pressure for each gas.
  50. Electrochemical series The half cell potential values are standard values and are represented as the standard reduction potential values as shown in the table at the end which is also called Electrochemical Series. Stronger oxidizing agent Weaker oxidizing agent F2(S) + 2 e— H202(aq) + 2 H+(aq) + 2e- C12Cg) + 2 e- + 14 H+(aq) + 6 02(g) + 4 H+(aq) + 4 e Br2(/) + 2 e- Ag+(aq) + e— Fe3+(aq) + e— 02(S) + 2 (aq) + 2 e 12(s) + 2 e— 02(S) + 2 H20(1) Cu2+(aq) + 2 e— Sn4+(aq) + 2 e- Pb2+(aq) + 2e- Ni2+(aq) + 2e- Zn2+(aq) + 2c- A13+(aq) + 3 e- Mg2+(aq) + 2 e— Li+(aq) + e— -...........-.> 2 F (aq) .-.......-.-..> 2 cr3+(aq) + 7 H20(/) .....................-> Ag(S) -...........-> Fe2+(aq) —.............-> H202(aq) ...........-...> Cu(s) ......-........> sn2+(aq) -........-..> H2(g) -.....-...-...> Pb(s) ...............-..> Ni(s) .................-> Fe(s) -......--> Zn(s) ..................> 112Cg) ..........-.-.> Al(s) -........-.-..> Mg(s) -......-......-.> Na(s) + 2 OH-(aq) 2.87 1.78 1.51 1.36 133 1.23 1.09 0.80 0.77 0.70 0.54 0.40 034 0.15 -0.13 -026 -0.40 -0.45 -0.76 -0.83 —1.66 -2.37 -2.71 -3.04 Wea ker reducing agent Stronger red ucing agent
  51. Cell potential or emf of a cell The difference between the electrode potentials of two half cells is called cell potential. It is known as electromotive force (EMF) of the cell if no current is drawn from the cell. cell cathode anode For this equation we take oxidation potential of anode and reduction potential of cathode. Since anode is put on left and cathode on right, it follows therefore, Left Right cell For a Daniel cell, therefore 2 = 0.34+0.76 = 1.10 v cell — Zn + cu
  52. Cell diagram or representation of a cell The following conventions or notations are applied for writing the cell diagram in accordance with IUPAC recommendations. The Daniel cell is represented as follows: zn(s) I Zn2+(C1) Il cu2+(C2) I cucs) Anode half cell is written on the left hand side while cathode half cell on 1. 1. 2. 3. right hand side. zn(s) I Zn2+ (aq) ; Anodic chamber cu2+ I Cucs) (aq) Cathodic chamber A single vertical line separates the metal from aqueous solution of its own ions. A double vertical line represents salt bridge The molar concentration (C) is placed in brackets after the formula of the corresponding ion.
  53. 4. 5. The value of e.m.f. of the cell is written on the extreme right of the cell. For example, zn(s) I Zn2+(1M) Il Cu2+(1M) I cucs) E.M.F = +1.1 V If an inert electrode like platinum is involved in the construction of the cell, it may be written along with the working electrode in bracket say for example, when a zinc anode Is connected to a hydrogen electrode. zn(s) I Zn2+(C1) Il H+(C2) I H2
  54. Salt bridge 1. 2. 3. Salt bridge is used to maintain the charge balance and to complete the circuit by facilitating the flow of ions through it. It contains a gel in which an inert electrolyte like Na2S04 or KN03 etc are mixed. Negative ions flow to the anode and positive ions flow to the cathode through the salt bridge and charge balance is maintained and cell keeps on functioning.
  55. Zinc Sol, VolimctctJAmme1cr soi 'bridge —???• soiz—-——-—- —???????— CillElUdi2 Coppcr Sol. Anode ???????#?• Cu —+ CII zn-cuso,-.zr.scu-cu
  56. Spontaneity of a reaction AG = - nFE Cell 1. For a spontaneous cell reaction AG should be negative and cell potential should be positive. 2. If we take standard value of cell potential in the above equation we will obtain standard value of AG as well. i.e. AGO - nFEO Cell
  57. Types of electrodes 1. Metal-Metal Ion electrodes A metal rod/plate is dipped in an electrolyte solution containing metal ions. There is a potential difference between these two phases and this electrode can act as a cathode or anode both. Anode: M Mn++ ne- Cathode: Mll+ + ne- M
  58. 2. Gas Electrodes Electrode gases like H2, C12 etc are used with their respective ions. For example, H2 gas is used with a dilute solution of HCI (H + ions). The metal should be inert so that it does not react with the acid. Anode: 2H++ ze- Cathode: 2H++ ze- The hydrogen electrode is also used as the standard to measure other electrode potentials. Its own potential is set to 0 V as a reference. When it is used as a reference the concentration of dil. HCI is taken as 1 M and the electrode is called "Standard Hydrogen Electrode (SHE)".
  59. ? ? ? ? ? ? ? t ? ? ? ? ? ? 1 0 0 ? With ?
  60. 3. Metal-Insoluble salt electrode We use salts of some metals which are sparingly soluble with the metal itself as electrodes. For example, if we use AgCl with Ag there is a potential gap between these two phases which can be identified in the following reaction: Agcl(s) + e- + cl- This electrode is made by dipping a silver rod in a solution containing AgCl(s) and Cl- ions.
  61. 4. Calomel Electrode Mercury is used with two other phases, one is a calomel paste (Hg2C12) and electrolyte containing Cl- ions. Cathode : ze- 2Hg(l) + zcl- (aq) Anode : This electrode is also used as another standard to measure other potentials. Its standard form is also called Standard Calomel Electrode (SCE).
  62. TO salt bridge pt. Wire Saturated KCt solution EEh'2Cin and Hg paste Bletx:uy Calomel electrode
  63. 5. Redox Electrode In these electrodes two different oxidation states of the same metal are used in the same half cell. For example, Fe2+ and Fe3+ are dissolved in the same container and an inert electrode of platinum is used for the electron transfer. Following reactions can take place: Anode: Fe2+ Fe3+ + e- Cathode: Fe3+ + e- Fe2+

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