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Page 1: Chapter 20. Electrochemistry - Wikispaces20... · AP Chemistry Chapter 20 Electrochemistry - 1 - Chapter 20. Electrochemistry . Sample Exercise 20.1 (p. 845) The nickel-cadmium (nicad)

AP Chemistry Chapter 20 Electrochemistry

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Chapter 20. Electrochemistry Sample Exercise 20.1 (p. 845) The nickel-cadmium (nicad) battery, a rechargeable “dry cell” used in battery-operated devices, uses the following redox reaction to generate electricity: Cd(s) + NiO2(s) + 2 H2O(l) Cd(OH)2(s) + Ni(OH)2(s) Identify the substances that are oxidized and reduced, and indicate which is the oxidizing agent and which is the reducing agent. Practice Exercise 20.1 Identify the oxidizing and reducing agents in the following oxidation-reduction equation: 2 H2O(l) + Al(s) + MnO4

-(aq) Al(OH)4

-(aq) + MnO2(s)

20.2 Balancing Oxidation-Reduction Equations Practice:

Write half-reactions for a reaction between Cu2+ and Zn metal.

Overall:

Oxidation:

Reduction:

Balancing Equations by the Method of Half-Reactions

1. Write the overall unbalanced reaction (if available). 2. (i) Identify oxidized and reduced substances

(ii) Write down the two incomplete half reactions. 3. Balance each half reaction: (a) First, balance elements other than H and O. (b) Then balance O atoms by adding water on side deficient in O atoms. (c) Then balance H by adding H+ for acidic solutions (or OH- for alkaline solutions). (d) Finish by balancing charge by adding electrons. 4. Multiply each half-reaction by the appropriate factors so that the number of electrons gained equals electrons lost. 5. Now add the reactions and simplify.

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AP Chemistry Chapter 20 Electrochemistry

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Sample Exercise 20.2 (p. 849) Complete and balance the following equation by the method of half-reactions: Cr2O7

2-(aq) + Cl-

(aq) Cr3+(aq) + Cl2(g) (acidic solution)

Practice Exercise 20.2 Complete and balance the following oxidation-reduction equations using the method of half-reactions. Both reactions occur in acidic solution.

a) Cu(s) + NO3-(aq) Cu2+

(aq) + NO2(g)

b) Mn2+

(aq) + NaBiO3(s) Bi3+(aq) + MnO4

-(aq)

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AP Chemistry Chapter 20 Electrochemistry

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Balancing Equations for Reactions Occurring in Basic Solution • The same method as above is used, but OH– is added to “neutralize” the H+ used. • The equation must again be simplified by canceling like terms on both sides of the equation. Sample Exercise 20.3 (p. 850) Complete and balance this equation for a redox reaction that takes place in basic solution: CN-

(aq) + MnO4-(aq) CNO-

(aq) + MnO2(s) (basic solution) Practice Exercise 20.3 Complete and balance the following equations for oxidation-reduction reactions that occur in basic solution:

a) NO2-(aq) + Al(s) NH3(aq) + Al(OH)4

-(aq)

b) Cr(OH)3(s) + ClO-

(aq) CrO42-

(aq) + Cl2(g)

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AP Chemistry Chapter 20 Electrochemistry

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20.3 Voltaic (Galvanic) Cells

Sample Exercise 20.4 (p. 853) The following oxidation-reduction reaction is spontaneous: Cr2O7

2-(aq) + 14 H+

(aq) + 6 I-(aq) 2 Cr3+

(aq) + 3 I2(s) + 7 H2O(l) A solution containing K2Cr2O7 and H2SO4 is poured into one beaker, and a solution of KI is poured into another. A salt bridge is used to join the beakers. A metallic conductor that will not react with either solution (such as platinum foil) is suspended in each solution, and the two conductors are connected with wires through a voltmeter or some other device to detect an electric current. The resultant voltaic cell generates an electric current. Indicate the reaction occurring at the anode, the reaction at the cathode, the direction of electron and ion migrations, and the signs of the electrodes. (Draw a diagram.) Practice Exercise 20.4 The two half-reactions in a voltaic cell are Zn(s) Zn2+

(aq) + 2 e- ClO3

-(aq) + 6 H+

(aq) + 6 e- Cl-(aq) + 3 H2O(l)

a) Indicate which reaction occurs at the anode and which at the cathode. b) Which electrode is consumed in the cell reaction? c) Which electrode is positive?

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AP Chemistry Chapter 20 Electrochemistry

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20.4 Cell EMF Sample Exercise 20.5 (p. 858) For the Zn-Cu2+ voltaic cell shown in Figure 20.5, we have Zn(s) + Cu2+

(aq, 1M) Zn2+(aq, 1M) + Cu(s) Eo

cell = 1.10 V Given that the standard reduction potential of Zn2+ to Zn is -0.76 V, calculate the Eo

red for the reduction of Cu2+ to Cu. Cu2+

(aq, 1M) + 2 e- Cu(s) (0.34 V) Practice Exercise 20.5 A voltaic cell is based on the following half-reactions: In+

(aq) In3+(aq) + 2 e-

Br2(l) + 2 e- 2 Br-(aq)

The standard emf for this cell is 1.46 V. Using the data in Table 20.1 or your AP Chem packet, calculate Eo

red for the reduction of In3+ to In+. (-0.40 V)

Sample Exercise 20.6 (p. 858) Using the standard reduction potentials listed in Table 20.1 or your AP Chem packet, calculate the standard emf for the voltaic cell described in Sample Exercise 20.4, which is based on the following reaction: Cr2O7

2-(aq) + 14 H+

(aq) + 6 I-(aq) 2 Cr3+

(aq) + 3 I2(s) + 7 H2O(l) (0.79 V) Practice Exercise 20.6 Using the data in Table 20.1 or your AP Chem packet, calculate the standard emf for a cell that employs the following overall cell reaction:

2 Al(s) + 3 I2(s) 2 Al3+(aq) + 6 I-

(aq) (+2.20 V)

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AP Chemistry Chapter 20 Electrochemistry

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Sample Exercise 20.7 (p. 859) A voltaic cell is based on the following two standard half-reactions:

Cd2+(aq) + 2 e- Cd(s)

Sn2+(aq) + 2 e- Sn(s)

By using the data in Appendix E or your AP Chem packet, determine

a) the half-reactions that occur at the cathode and the anode, and b) the standard cell potential (0.267 V)

Practice Exercise 20.7 A voltaic cell is based on a Co2+/Co half-cell and an AgCl/Ag half-cell. a) What reaction occurs at the anode? b) What is the standard cell potential? (0.499 V)

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AP Chemistry Chapter 20 Electrochemistry

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Strengths of Oxidizing and Reducing Agents

Sample Exercise 20.8 (p. 861) Using Table 20.1 or your AP Chem packet, rank the following ions in order of increasing strength as oxidizing agents: NO3

-(aq), Ag+

(aq), Cr2O72-

(aq). Practice Exercise 20.8 Using Table 20.1 or your AP Chem packet, rank the following species from the strongest to the weakest reducing agent: I-

(aq), Fe(s), Al(s).

20.5 Free Energy and Redox Reactions

• For any electrochemical process Eo = Eo

red(reduction process) – Eored(oxidation process).

Sample Exercise 20.9 (p. 862) Using standard reduction potentials (Table 20.1 or AP Chem packet), determine whether the following reactions are spontaneous under standard conditions:

a) Cu(s) + 2 H+(aq) Cu2+

(aq) + H2(g) b) Cl2(g) + 2 I-

(aq) 2 Cl-(aq) + I2(s)

Practice Exercise 20.9 Using standard reduction potentials (Appendix E or AP Chem packet), determine whether the following reactions are spontaneous under standard conditions:

a) I2(s) + 5 Cu2+(aq) + 6 H2O(l) 2 IO3

-(aq) + 5 Cu(s) + H+

(aq) b) Hg2+

(aq) + 2 I-(aq) Hg(l) + I2(s)

c) H2SO3(aq) + 2 Mn(s) + 4 H+(aq) S(s) + 2 Mn2+

(aq) + 3 H2O(l)

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AP Chemistry Chapter 20 Electrochemistry

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EMF and ∆G Sample Exercise 20.10 (p. 864)

a) Use the standard reduction potentials listed in Table 20.1 or your AP Chem packet to calculate the standard free-energy change, ∆Go, and the equilibrium constant, K, at 298 K for the following reaction:

4 Ag(s) + O2(g) + 4 H+(aq) 4 Ag+

(aq) + 2 H2O(l) (∆Go = -170 kJ/mol, K = 9 x 1029)

b) Suppose the reaction in part (a) were written as 2 Ag(s) + ½ O2(g) + 2 H+

(aq) 2 Ag+(aq) + H2O(l)

What are the values of Eo, ∆Go and K when the reaction is written in this way? (Eo = +0.43 V, ∆Go = - 83 kJ/mol, K = 4 x 1014)

Practice Exercise 20.10 For the reaction: 3 Ni2+

(aq) + 2 Cr(OH)3(s) + 10 OH-(aq) 3 Ni(s) + 2 CrO4

2-(aq) + 8 H2O(l)

a) What is the value of n for this reaction? (6) b) Use the data in Appendix E or your AP Chem packet to calculate ∆Go for this reaction. (+87 kJ/mol)

c) Calculate K at T = 298 K (6 x 10-16)

20.6 Cell EMF Under Nonstandard Conditions Sample Exercise 20.11 (p. 866) Calculate the emf at 298 K generated by the cell described in Sample Exercise 20.4 when [Cr2O7

2-] = 2.0 M, [H+] = 1.0 M, [I-] = 1.0 M, and [Cr3+] = 1.0 x 10-5 M. Cr2O7

2-(aq) + 14 H+

(aq) + 6 I-(aq) 2 Cr3+

(aq) + 3 I2(s) + 7 H2O(l) (E = +0.89 V)

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Practice Exercise 20.11 Calculate the emf generated by the cell described in the practice exercise accompanying Sample Exercise 20.6 when [Al3+] = 4.0 x 10-3 M and [I-] = 0.010 M. (E = +2.36 V)

Sample Exercise 20.12 (p. 866) If the voltage of a Zn-H+ cell (like that in Figure 20.11) is 0.45 V at 25oC when [Zn2+] = 1.0 M and PH2 = 1.0 atm, what is the concentration of H+? ([H+] = 5.8 x 10-6 M) Practice Exercise 20.12 What is the pH of the solution in the cathode compartment of the cell pictured in Figure 20.11 when PH2 = 1.0 atm, [Zn2+] in the anode compartment is 0.10 M, and cell emf is 0.542 V? (pH = 4.23)

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AP Chemistry Chapter 20 Electrochemistry

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Concentration Cells: • A concentration cell is one whose emf is generated solely because of a concentration difference. Sample Exercise 20.13 (p. 870) A voltaic cell is constructed with two hydrogen electrodes. Electrode 1 has PH2 = 1.00 atm and an unknown concentration of H+

(aq). Electrode 2 is a standard hydrogen electrode ([H+] = 1.00 M, PH2 = 1.00 atm). At 298 K the measured cell voltage is 0.211 V, and the electrical current is observed to flow from electrode 1 through the external circuit to electrode 2. Calculate [H+] for the solution at electrode 1. What is its pH? (pH = 3.57) Practice Exercise 20.13 A concentration cell is constructed with two Zn(s)-Zn2+

(aq) half-cells. The first half-cell has [Zn2+] = 1.35 M, and the second half-cell has [Zn2+] = 3.75 x 10-4 M. a) Which half-cell is the anode of the cell? b) What is the emf of the cell? (0.105 V)

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 = Keq) • Thus, if we know the cell emf, we can calculate the equilibrium constant.

0592.0log

log0592.00

°=

−°=

nEK

Kn

E

eq

eq

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20.9 Electrolysis

Sample Exercise 20.14 (p. 878) Calculate the number of grams of aluminum produced in 1.00 hr by the electrolysis of molten AlCl3 if the electrical current is 10.0 A. (3.3.6 g Al) Practice Exercise 20.14

a) The half-reaction for formation of magnesium metal upon electrolysis of molten MgCl2 is Mg2+ + 2e- Mg. Calculate the mass of magnesium formed upon passage of a current of 60.0 A for a period of 4.00 x 103 s.

(30.2 g Mg)

b) How many seconds would be required to produce 50.0 g of Mg from MgCl2 if the current is 100.0 A?

(3.97 x 103 s)

Electrical Work Sample Exercise 20.15 (p. 880) Calculate the number of kilowatt-hours of electricity required to produce 1.0 x 103 kg of aluminum by electrolysis of Al3+ if the applied emf is 4.50 V. (1.34 x 104 kWh)

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Practice Exercise 20.15 Calculate the number of kilowatt-hours of electricity required to produce 1.00 kg of Mg by electrolysis of molten MgCl2 if the applied emf is 5.00 V. Assume that the process is 100% efficient. (11.0 kWh)

Electrochemistry Terminology Current = a flow of electrical charge through a conductor. The current at a particular cross-section is the rate of flow of the charge. The charge may be carried by electrons (in metals), ions (electrolyte solutions) or positive holes (semi-conductors). Measured in amperes. Coulomb is the central unit for electrochemistry, enables us to maneuver amongst charge, voltage, current, watts, electrons and mass. The Faraday relates the charge to the number of electrons. 1 V = 1 J/C 1 A = 1 C/s 1 W = 1 J/s 1 F = 96,500 C/mol e- Electrical work: energy units x time 1 watt = 1 J/s 1 watt = 1 V x 1 A # Watts = volts x amps or J x C C s Sample Integrative Exercise 20 (p. 880) The Ksp at 298 K for iron (II) fluoride is 2.4 x 10-6.

a) Write a half-reaction that gives the likely products of the two-electron reduction of FeF2(s) in water. b) Use the Ksp value and the standard reduction potential of Fe2+

(aq) to calculate the standard reduction potential for the half-reaction in part (a). (-0.606 V)

c) Rationalize the difference in the reduction potential for the half-reaction in part (a) with that for

Fe2+(aq).