Electrochemical Cells

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Barnard Powers
5 years ago
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1 Electrochemistry

2 Electrochemical Cells

3 The Voltaic Cell Electrochemical Cell = device that generates electricity through redox rxns 1 Voltaic (Galvanic) Cell An electrochemical cell that produces an electrical current from a spontaneous rxn 2 Electrolytic Cell An electrochemical cell that consumes an electrical current to drive a nonspontaneous rxn

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5 Cell Notation Ex/ Zn (s) Zn 2+ (aq) Cu 2+ (aq) Cu (aq)

6 Half-Cell Potentials E cell = E oxidation + E reduction E oxidation = -E reduction + E cell means the rxn is spontaneous when adding E values for the half-cells, do not multiply the half-cell E values, even if you need to multiply the half-rxns to balance the equation

7 Predicting the Spontaneous Direction

8 Predicting the Spontaneous Direction ² Use E cell to determine which way a redox rxn will proceed ² The half-rxn with the more (+) electrode potential attracts e - s more strongly and undergoes reduction Good oxidizing agents ² The half-rxn with the more (-) electrode potential repels e - s more strongly and undergoes oxidation Good reducing agents ² Any reduction reaction in the table is spontaneous when paired with the reverse of any of the rxns listed

9 Predicting the Spontaneous Direction Without calculating E cell, predict whether the rxns below are spontaneous or not: a. Fe (s) + Mg 2+ (aq) à Fe 2+ (aq) + Mg (s) b. Fe (s) + Pb 2+ (aq) à Fe 2+ (aq) + Pb (s) A solution contains both NaI and NaBr. Which oxidizing agent could you add to the solution to selectively oxidize I - (aq) but not Br- (aq) a. Cl 2 b. H 2 O 2 c. CuCl 2 d. HNO 3

10 Predicting the Spontaneous Direction Do metals dissolve in acids? Yes, most acids dissolve metals by the reduction of H + ions to hydrogen gas and the corresponding oxidation of the metal to its ion. Zn (s) + 2H + (aq) à Zn 2+ (aq) + H 2(g) Can we predict whether a metal will dissolve in acid? What s the oxidation ½ rxn and the reduction ½ rxn? Metals whose reduction half-rxns are listed below the reduction of H + to H 2 dissolve in acids, while metals listed above it do not.

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12 Predicting the Spontaneous Direction Which metal(s) dissolves in HNO 3 but not in HCl? Hint: the combination of half-cell potentials should be positive a. Ni b. Au c. Cu Solution: Cu

13 Cell Potential, Free Energy, and K

14 E cell, ΔG and K Since we have related E cell to whether or not a rxn is spontaneous or not, can we relate E cell to ΔG? What about to K?

15 E cell, ΔG and K For a spontaneous reaction (one that proceeds in the forward direction with the chemicals in their standard states) ΔG < 0 (negative) E > 0 (positive) K > 1

16 E cell, ΔG and K For a nonspontaneous reaction (one that proceeds in the reverse direction with the chemicals in their standard states) ΔG > 0 (positive) E < 0 (negative) K < 1

17 Terms to know: Recall Electric Current = the flow of electric charge Current (I) is often represented by the number of electrons passing through per second (Amperes) Voltage = the potential energy per electron

18 E cell, ΔG and K Are we able to convert charge to numbers of electrons? Why, yes! Faraday s constant (F) represents the charge in coulombs of 1 mol of electrons: F = 96,485 coulombs/mol e - s (on your reference sheet)

19 E cell, ΔG and K Mathematically, we can derive a relationship between ΔG and E cell : ΔG = RTlnK ΔG = nfe cell n is the number of electrons F = Faraday s Constant = 96,485 C/mol e

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21 E cell, ΔG and K Ex/ Use the tabulated electrode potential values to calculate ΔG for the reaction: I 2(s) + 2Br - (aq) à 2I - (aq) + Br 2(l). Is the rxn spontaneous? Steps: 1 Label and separate the rxn into ½ reactions 2 Find E cell using your table of values 3 The value of n corresponds to the number of e - s canceled in the half-rxns. 4 Plug into ΔG = nfe cell. Note: 1V = J/C

22 Example - Calculate ΔG for the reaction I 2(s) + 2 Br (aq) Br 2(l) + 2 I (aq) Given: Find: Concept Plan: Relationships: Solve: Answer: I 2(s) + 2 Br (aq) Br 2(l) + 2 I (aq) ΔG, (J) E ox, E red E cell ΔG E = E + E! cell! ox! red!! G = nfe cell ox:! 2 Br (aq)! Br 2(l) + 2 e E = 1.09 v ΔG = nfe cell red: I 2(l) ( + 2 e )( 2 I (aq) E = )( v )! C J ΔG = 2 mol e 96, C tot: I 2(l) + 2Br (aq) 2I mole (aq) + Br 2(l) E = 0.55 v ΔG! = J since ΔG is +, the reaction is not spontaneous in the forward direction under standard conditions Δ

23 E cell, ΔG and K Ex/ Use the tabulated electrode potential values to calculate ΔG for the reaction: 2Na (s) + 2H 2 O (l) à H 2(g) + 2OH - (aq) + 2Na + (aq). Is the rxn spontaneous? Steps: 1 Label and separate the rxn into ½ reactions 2 Find E cell using your table of values 3 The value of n corresponds to the number of e - s canceled in the half-rxns. 4 Plug into ΔG = nfe cell. Note: 1V = J/C

24 Note: At this point, we would normally look at nonstandard conditions and thus the Nernst Equation I am skipping all notes on the Nernst equation, since they removed all AP standards related to it with the new test. However, I DO recommend reviewing it before taking college chemistry. I am also skipping over batteries. They are good to review as example problems, but focus on the halfrxns more than the definitions

25 Electrolysis

26 Electrolysis Recall: an electrolytic cell requires electrical energy to drive a nonspontaneous reaction For example, the reaction of hydrogen with oxygen to form water is spontaneous and can be used to produce an electrical current in a fuel cell Conversely, by supplying an electrical current, we can cause the reverse reaction to occur, separating water into hydrogen and oxygen

27 Electrolysis Electrolysis literally means to split with electricity Electrolytic cells Nonspontaneous the revrse rxn can occur if we supply a voltage greater than E cell we calculate for the spontaneous direction Used to separate ores or plate out metals We still follow AN OX & RED CAT

28 Electrolytic vs. Voltaic Cells Spontaneous vs. Nonspontaneous A voltaic cell is separated into two half-cells to generate electricity; an electrolytic cell often occurs in a single container A voltaic cell is a battery, while an electrolytic cell needs a battery AN OX, RED CAT However... For electrolytic cells: EPA electrolytic positive anode

29 Electrochemical Cells in all electrochemical cells, oxidation occurs at the anode, reduction occurs at the cathode in voltaic cells, anode is the source of electrons and has a ( ) charge cathode draws electrons and has a (+) charge in electrolytic cells electrons are drawn off the anode, so it must have a place to release the electrons, the + terminal of the battery electrons are forced toward the anode, so it must have a source of electrons, the terminal of the battery

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31 electroplating In electroplating, the work piece is the cathode. Cations are reduced at cathode and plate to the surface of the work piece. The anode is made of the plate metal. The anode oxidizes and replaces the metal cations in the solution

32 Electrolysis Electrolysis is the process of using electricity to break a compound apart electrolysis is done in an electrolytic cell electrolytic cells can be used to separate elements from their compounds generate H 2 from water for fuel cells recover metals from their ores

33 Electrolysis of Water

34 Electrolysis Products

35 Electrolysis of Pure Compounds Pure salts must be in molten (liquid) state electrodes normally graphite cations are reduced at the cathode to metal element anions oxidized at anode to nonmetal element Ex/ Oxidation: 2Cl - (l) à Cl 2(g) + 2e - Reduction: 2Na + (l) + 2e - à 2Na (s) 2Na + (l) + 2Cl - (l) à Cl 2(g) + 2Na (s)

36 Electrolysis of NaCl(l)

37 Mixtures of Ions when more than one cation is present, the cation that is easiest to reduce will be reduced first at the cathode least negative or most positive E red when more than one anion is present, the anion that is easiest to oxidize will be oxidized first at the anode least negative or most positive E ox

38 Aqueous Solutions Aqueous solutions can complicate things because of the presence of water There is the possibility of the electrolysis of water itself

39 Summing Up If there is no water present and you have a pure molten ionic compound, the cation will be reduced, and the anion will be oxidized If water is present & you have an aqueous solution of the ionic compound, then water will be oxidized instead No group 1A or IIA metal will be reduced in an aqueous solution à water will be oxidized instead

40 Stoichiometry of Electrolysis

41 Stoichiometry of Electrolysis Recall: In an electrolytic cell, electrical current is used to drive a rxn We can look at e - s as a reactant and use stoichiometric relationships For e - s, we measure quantity as charge For example, given Cu 2+ (aq) + 2e - à Cu (s) We can see 2 mol e - : 1 mol Cu (s) We can also relate mol of e - to charge using Faraday s constant: F = 96,485 C / mol e -

42 Faraday s Law The amount of metal deposited during electrolysis is directly proportional to the charge on the cation, the current, and the length of time the cell runs charge that flows through the cell = current x time 1 A = 1 C/s

43 Example - Calculate the mass of Au that can be plated in 25 min using 5.5 A for the half-reaction Au 3+ (aq) + 3 e Au(s) Given: Concept Plan: Relationships: Solve: Find: 25 min = 5.6 g Au 3 mol e : 1 mol Au, current = 5.5 amps, time = 25 min mass Au, g t(s), amp charge (C) mol e mol Au g Au 60 s 1min 5.5 C 1s 5.5 C 1s 1mol e 96,485 C 1mol Au 3 mol e 1mol e 1mol Au g 96,485 C 3 mol e 1mol Au g 1mol Au Check: units are correct, answer is reasonable since 10 A running for 1 hr ~ 1/3 mol e

44 Try This: Faraday s Law How long (in minutes) must a current of 5.00 A be applied to a solution of Ag + to produce 10.5 g silver metal? Hint: this is a reversal of what was being asked in the example Solution: 31.3 min

Electrochemistry Pearson Education, Inc. Mr. Matthew Totaro Legacy High School AP Chemistry

Electrochemistry Pearson Education, Inc. Mr. Matthew Totaro Legacy High School AP Chemistry 2012 Pearson Education, Inc. Mr. Matthew Totaro Legacy High School AP Chemistry Electricity from Chemistry Many chemical reactions involve the transfer of electrons between atoms or ions electron transfer