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Electrochemistry Chapter 20 Electrochemistry Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten.

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Presentation on theme: "Electrochemistry Chapter 20 Electrochemistry Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten."— Presentation transcript:

1 Electrochemistry Chapter 20 Electrochemistry Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

2 Electrochemistry Electrochemical Reactions In electrochemical reactions, electrons are transferred from one species to another. Metals tend to lose electrons and are oxidized, non metals tend to gain electrons and are reduced.

3 Electrochemistry LEO GER Losing Losing Electrons is Electrons is Oxidation. Oxidation. Gaining Electrons is Reduction

4 Electrochemistry OIL RIG OxidationOxidation Is Is Loss. Loss. ReductionReduction IsIs Gain. Gain.

5 Electrochemistry REDOX REACTIONS Are reduction – oxidation reactions.Are reduction – oxidation reactions. Electrons are transferred.Electrons are transferred. When an atom is losing electrons its O.N. increases. It is being oxidized.When an atom is losing electrons its O.N. increases. It is being oxidized. When an atom gains electrons its O.N. decreases. It is being reducedWhen an atom gains electrons its O.N. decreases. It is being reduced

6 Electrochemistry Oxidation Numbers In order to keep track of what loses electrons and what gains them, we assign oxidation numbers.

7 Electrochemistry Assigning Oxidation Numbers 1.Elements in their elemental form have an ON= 0. 2.The oxidation number of a monatomic ion is its charge. 3.Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.  Oxygen has an oxidation number of −2, except in the peroxide ion in which it has an oxidation number of −1.  Hydrogen is +1 except in metal hydrides when is −1.  Fluorine always has an oxidation number of −1.

8 Electrochemistry Assigning Oxidation Numbers  The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions 4.The sum of the oxidation numbers in a polyatomic ion is the charge on the ion. 5.The sum of the oxidation numbers in a neutral compound is 0.

9 Electrochemistry Oxidation and Reduction A species is oxidized when it loses electrons.  Here, zinc loses two electrons to go from neutral zinc metal to the Zn 2+ ion.

10 Electrochemistry Oxidation and Reduction A species is reduced when it gains electrons.  Here, each of the H + gains an electron and they combine to form H 2.

11 Electrochemistry Oxidation and Reduction What is reduced is the oxidizing agent.  H + oxidizes Zn by taking electrons from it. What is oxidized is the reducing agent.  Zn reduces H + by giving it electrons.

12 Electrochemistry Zn added to HCl yields the spontaneous reaction Zn(s) + 2H + (aq)  Zn 2+ (aq) + H 2 (g). The oxidation number of Zn has increased from 0 to 2+. The oxidation number of H has reduced from 1+ to 0. Zn is oxidized to Zn 2+ while H + is reduced to H 2. H + causes Zn to be oxidized and is the oxidizing agent. Zn causes H + to be reduced and is the reducing agent. Note that the reducing agent is oxidized and the oxidizing agent is reduced. Oxidation-Reduction Reactions

13 Electrochemistry Law of conservation of mass: the amount of each element present at the beginning of the reaction must be present at the end. Conservation of charge: electrons are not lost in a chemical reaction. Half Reactions Half-reactions are a convenient way of separating oxidation and reduction reactions. Balancing Oxidation- Reduction Reactions

14 Electrochemistry Half Reactions The half-reactions for Sn 2+ (aq) + 2Fe 3+ (aq)  Sn 4+ (aq) + 2Fe 2+ (aq) are Sn 2+ (aq)  Sn 4+ (aq) +2e - 2Fe 3+ (aq) + 2e -  2Fe 2+ (aq) Oxidation: electrons are products. Reduction: electrons are reactants. Loss of Gain of Electrons isElectrons is OxidationReduction

15 Electrochemistry Balancing Equations by the Method of Half Reactions Consider the titration of an acidic solution of Na 2 C 2 O 4 (sodium oxalate, colorless) with KMnO 4 (deep purple). MnO 4 - is reduced to Mn 2+ (pale pink) while the C 2 O 4 2- is oxidized to CO 2. The equivalence point is given by the presence of a pale pink color. If more KMnO 4 is added, the solution turns purple due to the excess KMnO 4.

16 Electrochemistry What is the balanced chemical equation? 1. Write down the two half reactions. 2. Balance each half reaction: a. First with elements other than H and O. b. Then balance O by adding water. c. Then balance H by adding H +. d. If it is in basic solution, remove H + by adding OH - e. Finish by balancing charge by adding electrons. 3. Multiply each half reaction to make the number of electrons equal. 4. Add the reactions and simplify. 5. Check!

17 Electrochemistry Half-Reaction Method Consider the reaction between MnO 4 − and C 2 O 4 2− : MnO 4 − (aq) + C 2 O 4 2− (aq)  Mn 2+ (aq) + CO 2 (aq)

18 Electrochemistry Balancing Equations by the Method of Half Reactions Consider the titration of an acidic solution of Na 2 C 2 O 4 (sodium oxalate, colorless) with KMnO 4 (deep purple). MnO 4 - is reduced to Mn 2+ (pale pink) while the C 2 O 4 2- is oxidized to CO 2. The equivalence point is given by the presence of a pale pink color. If more KMnO 4 is added, the solution turns purple due to the excess KMnO 4.

19 Electrochemistry Half-Reaction Method First, we assign oxidation numbers. MnO 4 − + C 2 O 4 2-  Mn 2+ + CO Since the manganese goes from +7 to +2, it is reduced. Since the carbon goes from +3 to +4, it is oxidized.

20 Electrochemistry For KMnO 4 + Na 2 C 2 O 4 : The two incomplete half reactions are MnO 4 - (aq)  Mn 2+ (aq) C 2 O 4 2- (aq)  2CO 2 (g) 2. Adding water and H + yields 8H + + MnO 4 - (aq)  Mn 2+ (aq) + 4H 2 O There is a charge of 7+ on the left and 2+ on the right. Therefore, 5 electrons need to be added to the left: 5e - + 8H + + MnO 4 - (aq)  Mn 2+ (aq) + 4H 2 O

21 Electrochemistry In the oxalate reaction, there is a 2- charge on the left and a 0 charge on the right, so we need to add two electrons: C 2 O 4 2- (aq)  2CO 2 (g) + 2e - 3. To balance the 5 electrons for permanganate and 2 electrons for oxalate, we need 10 electrons for both. Multiplying gives: 10e H + + 2MnO 4 - (aq)  2Mn 2+ (aq) + 8H 2 O 5C 2 O 4 2- (aq)  10CO 2 (g) + 10e -

22 Electrochemistry 4. Adding gives: 16H + (aq) + 2MnO 4 - (aq) + 5C 2 O 4 2- (aq)  2Mn 2+ (aq) + 8H 2 O(l) + 10CO 2 (g) 5. Which is balanced!

23 Electrochemistry Balancing in Basic Solution If a reaction occurs in basic solution, one can balance it as if it occurred in acid. Once the equation is balanced, add OH − to each side to “neutralize” the H + in the equation and create water in its place. If this produces water on both sides, you might have to subtract water from each side.

24 Electrochemistry Examples – Balance the following oxidation- reduction reactions: 1.Cr (s) + NO 3 - (aq)  Cr 3+ (aq) + NO (g) (acidic) 2.Al (s) + MnO 4 - (aq)  Al 3+ (aq) + Mn 2+ (aq) (acidic) 3.PO 3 3- (aq) + Mn 4 - (aq)  PO 4 3- (aq) + MnO 2 (s) (basic) 4.H 2 CO (aq) + Ag(NH 3 ) 2 + (aq)  HCO 3 - (aq) + Ag (s) + NH 3 (aq)(basic)

25 Electrochemistry Section Voltaic Cells – Spontaneous reactions ELECTRODES – POLARITIES SALT BRIDGE DRIVING FORCE- EMF STANDARD REDUCTION POTENTIALS HW Q 23, 25, 31, 33 (a, b ), 35

26 Electrochemistry Voltaic Cells In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released.

27 Electrochemistry Voltaic Cells We can use that energy to do work if we make the electrons flow through an external device. We call such a setup a voltaic cell.

28 Electrochemistry Voltaic Cells A typical cell looks like this. The oxidation occurs at the anode. The reduction occurs at the cathode.

29 Electrochemistry Voltaic Cells Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop.

30 Electrochemistry Voltaic Cells Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced.  Cations move toward the cathode.  Anions move toward the anode.

31 Electrochemistry Voltaic Cells In the cell, then, electrons leave the anode and flow through the wire to the cathode. As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.

32 Electrochemistry Voltaic Cells As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode. The electrons are taken by the cation, and the neutral metal is deposited on the cathode.

33 Electrochemistry Electromotive Force (emf) Water only spontaneously flows one way in a waterfall. Likewise, electrons only spontaneously flow one way in a redox reaction—from higher to lower potential energy.

34 Electrochemistry Electromotive Force (emf) The potential difference between the anode and cathode in a cell is called the electromotive force (emf). It is also called the cell potential, and is designated E cell.

35 Electrochemistry Cell Potential Cell potential is measured in volts (V). One volt is the potential difference required to impart 1J of energy to a charge of 1 coulomb. (1 electron has a charge of 1.6 x Coulombs). The potential difference between the 2 electrode provides the driving force that pushes the electron through the external circuit.

36 Electrochemistry Cell Potential or Electromotive Force (emf) The “pull” or driving force on the electrons. ELECTROMOTIVE FORCE EMF CAUSES THE ELECTRON MOTION! 1 V = 1 JCJC

37 Electrochemistry Standard Reduction Potentials Reduction potentials for many electrodes have been measured and tabulated.

38 Electrochemistry Standard Hydrogen Electrode Their values are referenced to a standard hydrogen electrode (SHE). By definition, the reduction potential for hydrogen is 0 V: 2 H + (aq, 1M) + 2 e −  H 2 (g, 1 atm)

39 Electrochemistry Standard Cell Potentials The cell potential at standard conditions can be found through this equation: E cell  = E red (cathode) − E red (anode)  Because cell potential is based on the potential energy per unit of charge, it is an intensive property.

40 Electrochemistry For the reaction to be SPONTANEOUS E has to be positive

41 Electrochemistry Cell Potentials For the oxidation in this cell, For the reduction, E red = −0.76 V  E red = V 

42 Electrochemistry Cell Potentials E cell  = E red  (cathode) − E red  (anode) = V − (−0.76 V) = V

43 Electrochemistry Since E  red = V we conclude that the reduction of Zn 2+ in the presence of the SHE is not spontaneous. The oxidation of Zn with the SHE is spontaneous. Changing the stoichiometric coefficient does not affect E  red. Therefore, 2Zn 2+ (aq) + 4e -  2Zn(s), E  red = V. Reactions with E  red > 0 are spontaneous reductions relative to the SHE.

44 Electrochemistry Reactions with E  red < 0 are spontaneous oxidations relative to the SHE. The larger the difference between E  red values, the larger E  cell. In a voltaic (galvanic) cell (spontaneous) E  red (cathode) is more positive than E  red (anode). Recall

45 Electrochemistry PNEUMONICS LEO GER OIL RIG RED CAT AND ALWAYS THE SOURCE OF ELECTRONS IS THE NEGATIVE ELECTRODE!!!

46 Electrochemistry In a voltaic (galvanic) cell (spontaneous) E  red (cathode) is more positive than E  red (anode) since Or More generally, for any electrochemical process A positive E  indicates a spontaneous process (galvanic cell). A negative E  indicates a nonspontaneous process. Spontaneity of Redox Reactions

47 Electrochemistry Cell Potential Calculations To Calculate cell potential using Standard Reduction Potentials: 1. One reaction and its cell potential’s sign must be reversed--it must be chosen such that the overall cell potential is positive. 2. The half-reactions must often be multiplied by an integer to balance electrons--this is not done for the cell potentials.

48 Electrochemistry Cell Potential Calculations Continued Fe 3+ (aq) + Cu (s) ----> Cu 2+ (aq) + Fe 2+ (aq) Fe 3+ (aq) + e > Fe 2+ (aq) E o = 0.77 V Cu 2+ (aq) + 2 e > Cu (s) E o = 0.34 V Reaction # 2 must be reversed.

49 Electrochemistry Cell Potential Calculations Continued 2 (Fe 3+ (aq) + e > Fe 2+ (aq) ) E o = 0.77 V Cu (s) ----> Cu 2+ (aq) + 2 e - E o = V 2Fe 3+ (aq) + Cu (s) ----> Cu 2+ (aq) + 2Fe 2+ (aq) E o = 0.43 V

50 Electrochemistry Oxidizing and Reducing Agents The greater the difference between the two, the greater the voltage of the cell. REMEMBER TO REVERSE THE SIGN OF THE SPECIE THAT GETS OXIDIZED (THE ONE BELOW!!!!!)

51 Electrochemistry Section 4.5 – 4.6 Free energy and redox rx Cell EMF under nonstandard conditions Nerst Equation- Concentration cells Redox Applications - Batteries and fuel cells – Corrosion prevention HW 47, 51 a, 59, 63

52 Electrochemistry Oxidizing and Reducing Agents The strongest oxidizers have the most positive reduction potentials. O.A. PULL electrons The strongest reducers have the most negative reduction potentials. R.A. PUSH their electrons

53 Electrochemistry Examples – Sketch the cells containing the following reactions. Include E° cell, the direction of electron flow, direction of ion migration through salt bridge, and identify the anode and cathode: 1.Cr 3+ + Cl 2 (g)  Cr 2 O Cl - 2.Cu 2+ + Mg (s)  Mg 2+ + Cu (s) 3.IO Fe 2+  Fe 3+ + I 2 4.Zn (s) + Ag +  Zn 2+ + Ag

54 Electrochemistry Free Energy  G for a redox reaction can be found by using the equation  G = −nFE where n is the number of moles of electrons transferred, and F is a constant, the Faraday. 1 F = 96,485 C/mol = 96,485 J/V-mol

55 Electrochemistry Free Energy and Cell Potential Under standard conditions  G  =  nFE  n = number of moles of electrons F = Faraday = 96,485 coulombs per mole of electrons = 96,485 J/Vmol E = V

56 Electrochemistry Find the free energy for the reaction of Zn + Cu > Zn 2+ + Cu

57 Electrochemistry The Nernst Equation A voltaic cell is functional until E = 0 at which point equilibrium has been reached. The point at which E = 0 is determined by the concentrations of the species involved in the redox reaction. The Nernst equation relates emf to concentration using and noting that Effect of Concentration on Cell EMF

58 Electrochemistry The Nernst Equation This rearranges to give the Nernst equation: The Nernst equation can be simplified by collecting all the constants together using a temperature of 298 K: (Note that change from natural logarithm to base-10 log.) Remember that n is number of moles of electrons.

59 Electrochemistry Cell EMF and Chemical Equilibrium A system is at equilibrium when  G = 0. From the Nernst equation, at equilibrium and 298 K (E = 0 V and Q = K eq ):

60 Electrochemistry Examples 1.Calculate E° cell for the following reaction: Zn (s) + Cu 2+  Zn 2+ + Cu (s) 2.Calculate E cell if [Zn 2+ ] = 1.88 M and [Cu 2+ ] = M 3.Calculate E° cell for the following reaction: IO 3 - (aq) + Fe 2+ (aq)  Fe 3+ (aq) + I 2 (aq) 4.Calculate E cell if [IO 3 - ] = 0.20 M, [Fe 2+ ] = 0.65 M, [Fe 3+ ] = 1.0 M, and [I 2 ] = 0.75 M

61 Electrochemistry Concentration Cells Notice that the Nernst equation implies that a cell could be created that has the same substance at both electrodes. For such a cell,would be 0, but Q would not. E cell  Therefore, as long as the concentrations are different, E will not be 0.

62 Electrochemistry Concentration Cells We can use the Nernst equation to generate a cell that has an emf based solely on difference in concentration. One compartment will consist of a concentrated solution, while the other has a dilute solution. Example: 1.00 M Ni 2+ (aq) and 1.00  M Ni 2+ (aq). The cell tends to equalize the concentrations of Ni 2+ (aq) in each compartment. The concentrated solution has to reduce the amount of Ni 2+ (aq) (to Ni(s)), so must be the cathode.

63 Electrochemistry Concentration Cells

64 Electrochemistry PROBLEM Use the Nernst equation to determine the (emf) at 25oC of the cell. Zn(s)│Zn2+( M ║Cu2+(10.0M)│Cu(s). The standard emf of this cell is 1.10 V.

65 Electrochemistry Applications of Oxidation-Reduction Reactions

66 Electrochemistry Batteries A battery is a galvanic cell or, more commonly, a group of galvanic cells connected in series.

67 Electrochemistry Batteries

68 Electrochemistry A lead storage battery consists of a lead anode, lead dioxide cathode, and an electrolyte of 38% sulfuric acid.

69 Electrochemistry Lead Storage Battery Anode reaction: Pb (s) +H 2 SO 4(aq) ---> PbSO 4(aq) + 2H + (aq) + 2e - Cathode reaction: PbO 2(s) +H 2 SO 4(aq) + 2e - +2H + (aq) - >PbSO 4(aq) +2H 2 O (l) Overall reaction: Pb (s) + PbO 2(s) + 2H 2 SO 4(aq) -->PbSO 4(aq) + 2H 2 O (l)

70 Electrochemistry Common dry cell and its components.

71 Electrochemistry Alkaline Battery Anode: Zn cap: Zn(s)  Zn 2+ (aq) + 2e - Cathode: MnO 2, NH 4 Cl and C paste: 2NH 4 + (aq) + 2MnO 2 (s) + 2e -  Mn 2 O 3 (s) + 2NH 3 (aq) + 2H 2 O(l) The graphite rod in the center is an inert cathode. For an alkaline battery, NH 4 Cl is replaced with KOH.

72 Electrochemistry Anode: Zn powder mixed in a gel: Zn(s)  Zn 2+ (aq) + 2e - Cathode: reduction of MnO 2.

73 Electrochemistry

74 Fuel Cells Direct production of electricity from fuels occurs in a fuel cell. On Apollo moon flights, the H 2 -O 2 fuel cell was the primary source of electricity. Cathode: reduction of oxygen: 2H 2 O(l) + O 2 (g) + 4e -  4OH - (aq) Anode: 2H 2 (g) + 4OH - (aq)  4H 2 O(l) + 4e -

75 Electrochemistry Mercury battery used in calculators.

76 Electrochemistry Fuel Cells...galvanic cells for which the reactants are continuously supplied. 2H 2 (g) + O 2 (g)  2H 2 O(l) anode: 2H 2 + 4OH   4H 2 O + 4e  cathode: 4e  + O 2 + 2H 2 O  4OH 

77 Electrochemistry Hydrogen Fuel Cells

78 Electrochemistry Ion selective electrodes are glass electrodes that measures a change in potential when [H + ] varies. Used to measure pH.

79 Electrochemistry Corrosion of Iron Since E  red (Fe 2+ ) < E  red (O 2 ) iron can be oxidized by oxygen. Cathode: O 2 (g) + 4H + (aq) + 4e -  2H 2 O(l). Anode: Fe(s)  Fe 2+ (aq) + 2e -. Dissolved oxygen in water usually causes the oxidation of iron. Fe 2+ initially formed can be further oxidized to Fe 3+ which forms rust, Fe 2 O 3.xH 2 O(s). Corrosion

80 Electrochemistry Oxidation occurs at the site with the greatest concentration of O 2. Preventing Corrosion of Iron Corrosion can be prevented by coating the iron with paint or another metal. Galvanized iron is coated with a thin layer of zinc.

81 Electrochemistry

82 Zinc protects the iron since Zn is the anode and Fe the cathode: Zn 2+ (aq) +2e -  Zn(s), E  red = V Fe 2+ (aq) + 2e -  Fe(s), E  red = V With the above standard reduction potentials, Zn is easier to oxidize than Fe.

83 Electrochemistry

84 Preventing Corrosion of Iron To protect underground pipelines, a sacrificial anode is added. The water pipe is turned into the cathode and an active metal is used as the anode. Often, Mg is used as the sacrificial anode: Mg 2+ (aq) +2e -  Mg(s), E  red = V Fe 2+ (aq) + 2e -  Fe(s), E  red = V

85 Electrochemistry

86 Corrosion Some metals, such as copper, gold, silver and platinum, are relatively difficult to oxidize. These are often called noble metals. About 1/5 of all iron and steel produced each year is used to replace rusted metal.

87 Electrochemistry Self-protecting Metals Some metals such as aluminum, copper, and silver form a protective coating that keeps them from corroding further. The protective coating for iron and steel flakes away opening new layers of metal to corrosion.

88 Electrochemistry Prevention of Corrosion Coating--painting or applying oil to keep out oxygen and moisture. Galvanizing--dipping a metal in a more active metal -- galvanized steel bucket. Alloying -- mixing metals with iron to prevent corrosion -- stainless steel. Cathodic protection -- attaching a more active metal. Serves as sacrificial metal--used to protect ships, gas lines, and gas tanks.

89 Electrochemistry Redox Titrations Same as any other titration. the permanganate ion is used often because it is its own indicator. MnO 4 - is purple, Mn +2 is colorless. When reaction solution remains clear, MnO 4 - is gone. Chromate ion is also useful, but color change, orangish yellow to green, is harder to detect.

90 Electrochemistry Example The iron content of iron ore can be determined by titration with standard KMnO 4 solution. The iron ore is dissolved in excess HCl, and the iron reduced to Fe +2 ions. This solution is then titrated with KMnO 4 solution, producing Fe +3 and Mn +2 ions in acidic solution. If it requires mL of M KMnO 4 to titrate a solution made with g of iron ore, what percent of the ore was iron?


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