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1 Electrochemistry Chapter 18, 19-20. 2 Electrochemical processes are oxidation-reduction reactions in which: the energy released by a spontaneous reaction.

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Presentation on theme: "1 Electrochemistry Chapter 18, 19-20. 2 Electrochemical processes are oxidation-reduction reactions in which: the energy released by a spontaneous reaction."— Presentation transcript:

1 1 Electrochemistry Chapter 18, 19-20

2 2 Electrochemical processes are oxidation-reduction reactions in which: the energy released by a spontaneous reaction is converted to electricity or electrical energy is used to cause a nonspontaneous reaction to occur

3 3 Balancing Redox Equations 1.Write the unbalanced equation for the reaction ion ionic form. The oxidation of Fe 2+ to Fe 3+ by Cr 2 O 7 2- in acid solution? Fe 2+ + Cr 2 O 7 2- Fe 3+ + Cr 3+ 2.Separate the equation into two half-reactions. Oxidation: Cr 2 O 7 2- Cr Reduction: Fe 2+ Fe Balance the atoms other than O and H in each half-reaction. Cr 2 O Cr 3+ Redox Reactions ( 氧化還原反應 ): 電化學的反應程序都是氧化還原反應。需理解氧化還原程式的平衡

4 4 Balancing Redox Equations 4.For reactions in acid, add H 2 O to balance O atoms and H + to balance H atoms. Cr 2 O Cr H 2 O 14H + + Cr 2 O Cr H 2 O 5.Add electrons to one side of each half-reaction to balance the charges on the half-reaction. Fe 2+ Fe e - 6e H + + Cr 2 O Cr H 2 O 6.If necessary, equalize the number of electrons in the two half- reactions by multiplying the half-reactions by appropriate coefficients. 6Fe 2+ 6Fe e - 6e H + + Cr 2 O Cr H 2 O

5 5 Balancing Redox Equations 7.Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel. 6e H + + Cr 2 O Cr H 2 O 6Fe 2+ 6Fe e - Oxidation: Reduction: 14H + + Cr 2 O Fe 2+ 6Fe Cr H 2 O 8.Verify that the number of atoms and the charges are balanced. 14x1 – x 2 = 24 = 6 x x 3 9.For reactions in basic solutions, add OH - to both sides of the equation for every H + that appears in the final equation.

6 6 由化學能轉換成電能 : 伏打電池 (voltaic cell) Galvanic Cells spontaneous redox reaction anode oxidation cathode reduction

7 7 Galvanic Cells The difference in electrical potential between the anode and cathode is called: cell voltage electromotive force (emf) cell potential Cell Diagram Zn (s) + Cu 2+ (aq) Cu (s) + Zn 2+ (aq) [Cu 2+ ] = 1 M and [Zn 2+ ] = 1 M Zn (s) | Zn 2+ (1 M) || Cu 2+ (1 M) | Cu (s) anode cathode salt bridge phase boundary

8 8 Standard Reduction Potentials Standard reduction potential (E ° ) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm. E° = 0 V Standard hydrogen electrode (SHE) 2e - + 2H + (1 M) H 2 (1 atm) Reduction Reaction

9 9 Standard Reduction Potentials Zn (s) | Zn 2+ (1 M) || H + (1 M) | H 2 (1 atm) | Pt (s) 2e - + 2H + (1 M) H 2 (1 atm) Zn (s) Zn 2+ (1 M) + 2e - Anode (oxidation): Cathode (reduction): Zn (s) + 2H + (1 M) Zn 2+ + H 2 (1 atm)

10 10 E 0 = 0.76 V cell Standard emf (E ° ) cell 0.76 V = 0 - E Zn /Zn ° 2+ E Zn /Zn = V ° 2+ Zn 2+ (1 M) + 2e - Zn E ° = V E ° = E H /H - E Zn /Zn cell °° Standard Reduction Potentials E ° = E cathode - E anode cell °° Zn (s) | Zn 2+ (1 M) || H + (1 M) | H 2 (1 atm) | Pt (s)

11 11 Standard Reduction Potentials Pt (s) | H 2 (1 atm) | H + (1 M) || Cu 2+ (1 M) | Cu (s) 2e - + Cu 2+ (1 M) Cu (s) H 2 (1 atm) 2H + (1 M) + 2e - Anode (oxidation): Cathode (reduction): H 2 (1 atm) + Cu 2+ (1 M) Cu (s) + 2H + (1 M) E ° = E cathode - E anode cell °° E ° = 0.34 V cell E cell = E Cu /Cu – E H /H °°° 0.34 = E Cu /Cu - 0 ° 2+ E Cu /Cu = 0.34 V 2+ °

12 12 E ° is for the reaction as written The more positive E ° the greater the tendency for the substance to be reduced The half-cell reactions are reversible The sign of E ° changes when the reaction is reversed Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E °

13 13 What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 M Cd(NO 3 ) 2 solution and a Cr electrode in a 1.0 M Cr(NO 3 ) 3 solution? Cd 2+ (aq) + 2e - Cd (s) E ° = V Cr 3+ (aq) + 3e - Cr (s) E ° = V Cd is the stronger oxidizer Cd will oxidize Cr 2e - + Cd 2+ (1 M) Cd (s) Cr (s) Cr 3+ (1 M) + 3e - Anode (oxidation): Cathode (reduction): 2Cr (s) + 3Cd 2+ (1 M) 3Cd (s) + 2Cr 3+ (1 M) x 2 x 3 E ° = E cathode - E anode cell °° E ° = – (-0.74) cell E ° = 0.34 V cell

14 14 Spontaneity of Redox Reactions  G = -nFE cell  G ° = -nFE cell ° n = number of moles of electrons in reaction F = 96,500 J V mol = 96,500 C/mol  G ° = -RT ln K = -nFE cell ° E cell ° = RT nF ln K (8.314 J/K mol)(298 K) n (96,500 J/V mol) ln K = = V n ln K E cell ° = V n log K E cell °

15 15 Spontaneity of Redox Reactions  G ° = -RT ln K = -nFE cell °

16 16 Batteries Leclanché cell Dry cell Zn (s) Zn 2+ (aq) + 2e - Anode: Cathode: 2NH 4 (aq) + 2MnO 2 (s) + 2e - Mn 2 O 3 (s) + 2NH 3 (aq) + H 2 O (l) + Zn (s) + 2NH 4 (aq) + 2MnO 2 (s) Zn 2+ (aq) + 2NH 3 (aq) + H 2 O (l) + Mn 2 O 3 (s)

17 17 Batteries Zn(Hg) + 2OH - (aq) ZnO (s) + H 2 O (l) + 2e - Anode: Cathode: HgO (s) + H 2 O (l) + 2e - Hg (l) + 2OH - (aq) Zn(Hg) + HgO (s) ZnO (s) + Hg (l) Mercury Battery

18 18 Batteries Anode: Cathode: Lead storage battery PbO 2 (s) + 4H + (aq) + SO 2- (aq) + 2e - PbSO 4 (s) + 2H 2 O (l) 4 Pb (s) + SO 2- (aq) PbSO 4 (s) + 2e - 4 Pb (s) + PbO 2 (s) + 4H + (aq) + 2SO 2- (aq) 2PbSO 4 (s) + 2H 2 O (l) 4

19 19 Batteries Solid State Lithium Battery

20 20 Batteries A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning Anode: Cathode: O 2 (g) + 2H 2 O (l) + 4e - 4OH - (aq) 2H 2 (g) + 4OH - (aq) 4H 2 O (l) + 4e - 2H 2 (g) + O 2 (g) 2H 2 O (l)

21 21 Corrosion Corrosion is the term usually applied to the deterioration of metals by an electrochemical process.

22 22 Cathodic Protection of an Iron Storage Tank

23 23 Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur. Electrolysis of molten NaCl

24 24 Electrolysis of Water

25 25 Electrolysis and Mass Changes charge (C) = current (A) x time (s) 1 mol e - = 96,500 C

26 26 How much Ca will be produced in an electrolytic cell of molten CaCl 2 if a current of A is passed through the cell for 1.5 hours? Anode: Cathode: Ca 2+ (l) + 2e - Ca (s) 2Cl - (l) Cl 2 (g) + 2e - Ca 2+ (l) + 2Cl - (l) Ca (s) + Cl 2 (g) 2 mole e - = 1 mole Ca mol Ca = C s x 1.5 hr x 3600 s hr96,500 C 1 mol e - x 2 mol e - 1 mol Ca x = mol Ca = 0.50 g Ca


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