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Chapter 18 Electrochemistry. Chapter 18 Table of Contents Copyright © Cengage Learning. All rights reserved 2 18.1Balancing Oxidation–Reduction Equations.

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Presentation on theme: "Chapter 18 Electrochemistry. Chapter 18 Table of Contents Copyright © Cengage Learning. All rights reserved 2 18.1Balancing Oxidation–Reduction Equations."— Presentation transcript:

1 Chapter 18 Electrochemistry

2 Chapter 18 Table of Contents Copyright © Cengage Learning. All rights reserved 2 18.1Balancing Oxidation–Reduction Equations 18.2 Galvanic Cells 18.3 Standard Reduction Potentials 18.4Cell Potential, Electrical Work, and Free Energy 18.5Dependence of Cell Potential on Concentration 18.6Batteries 18.7Corrosion 18.8Electrolysis 18.9Commercial Eletrolytic Processes

3 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 3 Review of Terms Oxidation – reduction (redox) reaction – involves a transfer of electrons from the reducing agent to the oxidizing agent Oxidation – loss of electrons Reduction – gain of electrons Reducing agent – electron donor Oxidizing agent – electron acceptor

4 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 4 Half–Reactions The overall reaction is split into two half – reactions, one involving oxidation and one reduction. 8H + + MnO 4  + 5Fe 2+  Mn 2+ + 5Fe 3+ + 4H 2 O Reduction: 8H + + MnO 4  + 5e   Mn 2+ + 4H 2 O Oxidation: 5Fe 2+  5Fe 3+ + 5e 

5 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 5 1.Write separate equations for the oxidation and reduction half – reactions. 2.For each half – reaction: A.Balance all the elements except H and O. B.Balance O using H 2 O. C.Balance H using H +. D.Balance the charge using electrons. The Half–Reaction Method for Balancing Equations for Oxidation– Reduction Reactions Occurring in Acidic Solution

6 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 6 3.If necessary, multiply one or both balanced half – reactions by an integer to equalize the number of electrons transferred in the two half – reactions. 4.Add the half – reactions, and cancel identical species. 5.Check that the elements and charges are balanced. The Half–Reaction Method for Balancing Equations for Oxidation– Reduction Reactions Occurring in Acidic Solution

7 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 7 The Half–Reaction Method for Balancing Equations for Oxidation– Reduction Reactions Occurring in Acidic Solution

8 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 8 Cr 2 O 7 2- (aq) + SO 3 2- (aq)  Cr 3+ (aq) + SO 4 2- (aq) How can we balance this equation? First Steps:  Separate into half-reactions.  Balance elements except H and O.

9 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 9 Cr 2 O 7 2- (aq)  2Cr 3+ (aq) SO 3 2- (aq)  SO 4 2- (aq) How many electrons are involved in each half reaction? Method of Half Reactions

10 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 10 6e - + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) SO 3 2- (aq)  + SO 4 2- (aq) + 2e - How can we balance the oxygen atoms? Method of Half Reactions (continued)

11 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 11 6e - + Cr 2 O 7 2- (aq)  Cr 3+ (aq) + 7H 2 O H 2 O +SO 3 2- (aq)  + SO 4 2- (aq) + 2e - How can we balance the hydrogen atoms? Method of Half Reactions (continued)

12 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 12 This reaction occurs in an acidic solution. 14H + + 6e - + Cr 2 O 7 2-  2Cr 3+ + 7H 2 O H 2 O +SO 3 2-  SO 4 2- + 2e - + 2H + How can we balance the electrons? Method of Half Reactions (continued)

13 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 13 14H + + 6e - + Cr 2 O 7 2-  2Cr 3+ + 7H 2 O 3[H 2 O +SO 3 2-  SO 4 2- + 2e - + 2H + ] Final Balanced Equation: Cr 2 O 7 2- + 3SO 3 2- + 8H +  2Cr 3+ + 3SO 4 2- + 4H 2 O Method of Half Reactions (continued)

14 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 14 Exercise Balance the following oxidation – reduction reaction that occurs in acidic solution. Br – (aq) + MnO 4 – (aq)  Br 2 (l)+ Mn 2+ (aq) 10Br – (aq) + 16H + (aq) + 2MnO 4 – (aq)  5Br 2 (l)+ 2Mn 2+ (aq) + 8H 2 O(l)

15 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 15 1.Use the half–reaction method as specified for acidic solutions to obtain the final balanced equation as if H + ions were present. 2.To both sides of the equation, add a number of OH – ions that is equal to the number of H + ions. (We want to eliminate H + by forming H 2 O.) The Half–Reaction Method for Balancing Equations for Oxidation– Reduction Reactions Occurring in Basic Solution

16 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 16 3.Form H 2 O on the side containing both H + and OH – ions, and eliminate the number of H 2 O molecules that appear on both sides of the equation. 4.Check that elements and charges are balanced. The Half–Reaction Method for Balancing Equations for Oxidation– Reduction Reactions Occurring in Basic Solution

17 Section 18.1 Balancing Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 17 The Half–Reaction Method for Balancing Equations for Oxidation– Reduction Reactions Occurring in Basic Solution

18 Section 18.2 Atomic MassesGalvanic Cells Return to TOC Copyright © Cengage Learning. All rights reserved 18 Galvanic Cell Device in which chemical energy is changed to electrical energy. Uses a spontaneous redox reaction to produce a current that can be used to do work.

19 Section 18.2 Atomic MassesGalvanic Cells Return to TOC Copyright © Cengage Learning. All rights reserved 19 A Galvanic Cell

20 Section 18.2 Atomic MassesGalvanic Cells Return to TOC Copyright © Cengage Learning. All rights reserved 20 Galvanic Cell Oxidation occurs at the anode. Reduction occurs at the cathode. Salt bridge or porous disk – devices that allow ions to flow without extensive mixing of the solutions.  Salt bridge – contains a strong electrolyte held in a Jello–like matrix.  Porous disk – contains tiny passages that allow hindered flow of ions.

21 Section 18.2 Atomic MassesGalvanic Cells Return to TOC Copyright © Cengage Learning. All rights reserved 21 Voltaic Cell: Cathode Reaction

22 Section 18.2 Atomic MassesGalvanic Cells Return to TOC Copyright © Cengage Learning. All rights reserved 22 Voltaic Cell: Anode Reaction

23 Section 18.2 Atomic MassesGalvanic Cells Return to TOC Copyright © Cengage Learning. All rights reserved 23 Cell Potential A galvanic cell consists of an oxidizing agent in one compartment that pulls electrons through a wire from a reducing agent in the other compartment. The “pull”, or driving force, on the electrons is called the cell potential ( ), or the electromotive force (emf) of the cell.  Unit of electrical potential is the volt (V).  1 joule of work per coulomb of charge transferred.

24 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 24 Electrochemical Half-Reactions in a Galvanic Cell

25 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 25 Galvanic Cell All half-reactions are given as reduction processes in standard tables.  Table 18.1  1 M, 1atm, 25°C When a half-reaction is reversed, the sign of E ° is reversed. When a half-reaction is multiplied by an integer, E ° remains the same. A galvanic cell runs spontaneously in the direction that gives a positive value for E ° cell.

26 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 26 Example: Fe 3+ (aq) + Cu(s) → Cu 2+ (aq) + Fe 2+ (aq) Half-Reactions:  Fe 3+ + e – → Fe 2+ E ° = 0.77 V  Cu 2+ + 2e – → Cu E ° = 0.34 V To balance the cell reaction and calculate the cell potential, we must reverse reaction 2.  Cu → Cu 2+ + 2e – – E ° = – 0.34 V Each Cu atom produces two electrons but each Fe 3+ ion accepts only one electron, therefore reaction 1 must be multiplied by 2.  2Fe 3+ + 2e – → 2Fe 2+ E ° = 0.77 V

27 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 27 Overall Balanced Cell Reaction 2Fe 3+ + 2e – → 2Fe 2+ E ° = 0.77 V (cathode) Cu → Cu 2+ + 2e – – E ° = – 0.34 V (anode) Balanced Cell Reaction: Cu + 2Fe 3+ → Cu 2+ + 2Fe 2+ Cell Potential: E ° cell = E °(cathode) – E °(anode) E ° cell = 0.77 V – 0.34 V = 0.43 V

28 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 28 Concept Check Order the following from strongest to weakest oxidizing agent and justify. Of those you cannot order, explain why. FeNaF - Na + Cl 2

29 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 29 Line Notation Used to describe electrochemical cells. Anode components are listed on the left. Cathode components are listed on the right. Separated by double vertical lines. The concentration of aqueous solutions should be specified in the notation when known. Example: Mg(s)|Mg 2+ (aq)||Al 3+ (aq)|Al(s)  Mg → Mg 2+ + 2e – (anode)  Al 3+ + 3e – → Al(cathode)

30 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 30 Description of a Galvanic Cell The cell potential is always positive for a galvanic cell where E ° cell = E °(cathode)–E °(anode) and the balanced cell reaction. The direction of electron flow, obtained by inspecting the half–reactions and using the direction that gives a positive E ° cell.

31 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 31 Description of a Galvanic Cell Designation of the anode and cathode. The nature of each electrode and the ions present in each compartment. A chemically inert conductor is required if none of the substances participating in the half–reaction is a conducting solid.

32 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 32 Concept Check Sketch a cell using the following solutions and electrodes. Include:  The potential of the cell  The direction of electron flow  Labels on the anode and the cathode a)Ag electrode in 1.0 M Ag + (aq) and Cu electrode in 1.0 M Cu 2+ (aq)

33 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 33 Concept Check Sketch a cell using the following solutions and electrodes. Include:  The potential of the cell  The direction of electron flow  Labels on the anode and the cathode b)Zn electrode in 1.0 M Zn 2+ (aq) and Cu electrode in 1.0 M Cu 2+ (aq)

34 Section 18.3 The MoleStandard Reduction Potentials Return to TOC Copyright © Cengage Learning. All rights reserved 34 Concept Check Consider the cell from part b. What would happen to the potential if you increase the [Cu 2+ ]? Explain. The cell potential should increase.

35 Section 18.4 Cell Potential, Electrical Work, and Free Energy Return to TOC Copyright © Cengage Learning. All rights reserved 35 Work Work is never the maximum possible if any current is flowing. In any real, spontaneous process some energy is always wasted – the actual work realized is always less than the calculated maximum.

36 Section 18.4 Cell Potential, Electrical Work, and Free Energy Return to TOC Copyright © Cengage Learning. All rights reserved 36 Maximum Cell Potential Directly related to the free energy difference between the reactants and the products in the cell.  ΔG° = –nFE °  F = 96,485 C/mol e –

37 Section 18.5 Dependence of Cell Potential on Concentration Return to TOC Copyright © Cengage Learning. All rights reserved 37 A Concentration Cell

38 Section 18.5 Dependence of Cell Potential on Concentration Return to TOC Copyright © Cengage Learning. All rights reserved 38 Nernst Equation At 25°C: or (at equilibrium)

39 Section 18.5 Dependence of Cell Potential on Concentration Return to TOC Copyright © Cengage Learning. All rights reserved 39 Concept Check Explain the difference between E and E °. When is E equal to zero? W hen the cell is in equilibrium ("dead" battery). When is E ° equal to zero? E  is equal to zero for a concentration cell.

40 Section 18.5 Dependence of Cell Potential on Concentration Return to TOC Copyright © Cengage Learning. All rights reserved 40 Exercise A concentration cell is constructed using two nickel electrodes with Ni 2+ concentrations of 1.0 M and 1.00 x 10 -4 M in the two half-cells. Calculate the potential of this cell at 25°C. 0.118 V

41 Section 18.5 Dependence of Cell Potential on Concentration Return to TOC Copyright © Cengage Learning. All rights reserved 41 Concept Check You make a galvanic cell at 25°C containing:  A nickel electrode in 1.0 M Ni 2+ (aq)  A silver electrode in 1.0 M Ag + (aq) Sketch this cell, labeling the anode and cathode, showing the direction of the electron flow, and calculate the cell potential. 1.03 V

42 Section 18.6 Batteries Return to TOC Copyright © Cengage Learning. All rights reserved 42 One of the Six Cells in a 12–V Lead Storage Battery

43 Section 18.6 Batteries Return to TOC Copyright © Cengage Learning. All rights reserved 43 A Common Dry Cell Battery

44 Section 18.6 Batteries Return to TOC Copyright © Cengage Learning. All rights reserved 44 A Mercury Battery

45 Section 18.6 Batteries Return to TOC Copyright © Cengage Learning. All rights reserved 45 Schematic of the Hydrogen-Oxygen Fuel Cell

46 Section 18.7 Corrosion Return to TOC Copyright © Cengage Learning. All rights reserved 46 Process of returning metals to their natural state – the ores from which they were originally obtained. Involves oxidation of the metal.

47 Section 18.7 Corrosion Return to TOC Copyright © Cengage Learning. All rights reserved 47 The Electrochemical Corrosion of Iron

48 Section 18.7 Corrosion Return to TOC Copyright © Cengage Learning. All rights reserved 48 Corrosion Prevention Application of a coating (like paint or metal plating)  Galvanizing Alloying Cathodic Protection  Protects steel in buried fuel tanks and pipelines.

49 Section 18.7 Corrosion Return to TOC Copyright © Cengage Learning. All rights reserved 49 Cathodic Protection

50 Section 18.8 Electrolysis Return to TOC Copyright © Cengage Learning. All rights reserved 50 Forcing a current through a cell to produce a chemical change for which the cell potential is negative.

51 Section 18.8 Electrolysis Return to TOC Copyright © Cengage Learning. All rights reserved 51 Stoichiometry of Electrolysis How much chemical change occurs with the flow of a given current for a specified time? current and time  quantity of charge  moles of electrons  moles of analyte  grams of analyte

52 Section 18.8 Electrolysis Return to TOC Copyright © Cengage Learning. All rights reserved 52 Stoichiometry of Electrolysis current and time  quantity of charge Coulombs of charge = amps (C/s) × seconds (s) quantity of charge  moles of electrons

53 Section 18.8 Electrolysis Return to TOC Copyright © Cengage Learning. All rights reserved 53 Concept Check An unknown metal (M) is electrolyzed. It took 52.8 sec for a current of 2.00 amp to plate 0.0719 g of the metal from a solution containing M(NO 3 ) 3. What is the metal? gold (Au)

54 Section 18.8 Electrolysis Return to TOC Copyright © Cengage Learning. All rights reserved 54 Concept Check Consider a solution containing 0.10 M of each of the following: Pb 2+, Cu 2+, Sn 2+, Ni 2+, and Zn 2+. Predict the order in which the metals plate out as the voltage is turned up from zero. Cu 2+, Pb 2+, Sn 2+, Ni 2+, Zn 2+ Do the metals form on the cathode or the anode? Explain.

55 Section 18.9 Commercial Electrolytic Processes Return to TOC Copyright © Cengage Learning. All rights reserved 55 Production of aluminum Purification of metals Metal plating Electrolysis of sodium chloride Production of chlorine and sodium hydroxide

56 Section 18.9 Commercial Electrolytic Processes Return to TOC Copyright © Cengage Learning. All rights reserved 56 Producing Aluminum by the Hall-Heroult Process

57 Section 18.9 Commercial Electrolytic Processes Return to TOC Copyright © Cengage Learning. All rights reserved 57 Electroplating a Spoon

58 Section 18.9 Commercial Electrolytic Processes Return to TOC Copyright © Cengage Learning. All rights reserved 58 The Downs Cell for the Electrolysis of Molten Sodium Chloride

59 Section 18.9 Commercial Electrolytic Processes Return to TOC Copyright © Cengage Learning. All rights reserved 59 The Mercury Cell for Production of Chlorine and Sodium Hydroxide


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