Ch. 18 Electrochemistry

1. Ch. 18 Electrochemistry Dr. Namphol Sinkaset Chem 201: General Chemistry II

2. I. Chapter Outline • Introduction • Balancing Redox Reactions • Galvanic Cells • Standard Reduction Potentials • E°cell, ΔG°, and K • Batteries • Electrolysis

3. I. Introduction • Reduction-oxidation (redox) reactions are a huge branch of chemistry. • Basically, they are reactions in which electrons are in motion. • Electrons in motion = electricity. • Concepts in this chapter can be used to explain fuel cells, batteries, electroplating, etc.

4. I. Redox Reactions • Recall that reduction and oxidation are coupled processes. • Reduction is the gain of electrons. • Oxidation is the loss of electrons. • We identify what is oxidized and what is reduced by examining oxidation numbers.

5. I. Assigning Oxidation Numbers • Atoms in elemental form have O.N. = 0. • Charge on a monatomic ion equals its O.N. • The sum of all O.N. must equal the total charge. • For Group 1, O.N. = +1. • For Group 2, O.N. = +2. • For H, O.N. = +1 w/ nonmetals, -1 w/ metals and B. • For F, O.N. = -1. • For O, O.N. = -1 in peroxides and -2 in all others. • For Group 17, typically O.N. = -1.

6. I. Sample Redox Reaction • Ca(s) is oxidized; it is the reducing agent. • H2O(l) is reduced; it is the oxidizing agent.

7. II. Balancing Redox Reactions • Balancing a redox reaction is more involved because mass and charge must both be balanced. • A common way to balance redox reactions is the “half-cell method.” • In this method, we split the reaction into 2 half reactions: oxidation and reduction. • Each half is balanced separately, and then the two are recombined.

8. II. The Half-Cell Method • Identify what is being oxidized and reduced. • Write each half-cell reaction. • Balance each half-cell w/ respect to mass: • Balance elements other than O and H. • Balance O by adding H2O. • Balance H by adding H+. • Balance each half-cell w/ respect to charge by adding e-’s. • Make the # of e-’s in each half-cell equal by multiplying by an appropriate factor. • Add the two half-cells together. • Add required # of OH- to each side and simplify if redox reaction takes place in basic solution.

9. II. Sample Problem • Balance the following reaction in acidic solution: ClO-(aq) + Cr(OH)4-(aq) CrO42-(aq) + Cl-(aq).

10. II. Sample Problem • Balance the following reaction in basic solution: MnO4-(aq) + Br-(aq) MnO2(s) + BrO3-(aq)

11. III. Generating Electricity • As stated earlier, electrical current is simply charge in motion. • Since an e- is charged, movement of electrons generates electrical current. • Some redox reaction occur spontaneously. • e.g. Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

12. III. Zn/Cu2+ Redox • Zn(s) is oxidized to Zn2+ which goes into solution. • Cu2+ is reduced to Cu(s) which “plates out.” • 2 e-’s are transferred from Zn to Cu2+.

13. III. Making e-’s Travel • In this arrangement, we don’t really get electrons flowing, more of a transfer. • If we separate the components, we can force electrons to travel externally, and even do work for us. • This kind of setup is called an electrochemical cell.

14. III. Electrochemical Cell

15. III. Electrochemical Cells • Voltaic (or galvanic) cells produce electrical current from a spontaneous redox reaction. • Important aspects: • Each part is called a half-cell. • Oxidation occurs at the anode. Electrons flow away from the anode. • Reduction occurs at the cathode. Electrons come to the cathode. • A salt bridge is necessary to prevent charge build up and complete the circuit.

16. III. Current and Potential • Two important aspects about electrons in motion are current and potential. • Electrical current is measured in amperes (A), which has units of coulombs (measure of charge) per second, C/s. 1 A = 1 C/s. • The current is driven by a potential energy difference called the potential difference. • Potential difference is a measure of the difference in PE per unit of charge. • The SI unit is the volt (V). 1 V = 1 J/C.

17. III. Cell Potentials • The potential difference is the force that drives the movement of e-’s, so it’s also called the electromotive force (emf). • In a voltaic cell, the potential difference between the cathode and anode is called the cell potential (Ecell) or cell emf.

18. III. Standard Cell Potentials • If the cell is under standard conditions, the cell potential is the standard cell potential (E°cell) or standard emf. • For the Zn(s)/Cu2+(aq) cell, E°cell = 1.10 V. • The cell potential is a measure of the tendency of the redox reaction to occur spontaneously.

19. III. Electrochemical Cell Notation • Instead of drawing an electrochemical cell, line notation can be used. • The Zn(s)/Cu2+(aq) can be represented as Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s) • Oxidation is always written first, followed by the reduction. • Double line represents the salt bridge. • Different phases are separated by single lines. • Multiple substances in solution are separated by commas.

20. III. Sample Cell Notation Fe(s)|Fe2+(aq)||MnO4-(aq), H+(aq), Mn2+(aq)|Pt(s)

21. III. Sample Problem • Draw and completely label the electrochemical cell represented by the line notation shown below. Additionally, write the overall balanced equation for the redox reaction. Sn(s)|Sn2+(aq)||NO(g)|NO3-(aq), H+(aq)|Pt(s)

22. IV. Calculating Cell Potentials • In an electrochemical cell, we can think of each half-cell as having its own potential. • Thus, E°cell is a sum of both half-cell potentials. • The half-cell with the higher potential will occur in the forward direction, forcing the other half-cell to occur in the reverse direction.

23. IV. Half-Cell Potentials • We can’t create a cell that has only reduction or only oxidation; thus, we can’t measure an absolute value of a half-cell. • We have to assign a particular half-cell a value of 0.00 V and measure everything else relative to that.

24. IV. Standard Hydrogen Electrode • The standard hydrogen electrode (SHE) is normally assigned a potential of 0.00 V. • We attach different half-cells to SHE, and whatever potential comes out is assigned to the half-cell.

25. IV. Measuring a Standard Reduction Potential

26. IV. Standard Reduction Potentials

27. IV. Important Points • The reduction potential of SHE is 0.00 V. • Half-cells with greater tendency to undergo reduction than SHE have a positive reduction potential. • Half-cells with lesser tendency to undergo reduction than SHE have a negative reduction potential. • Substances at the top have strong tendency to be reduced; they are strong oxidizing agents.

28. IV. Important Points • Substances at the bottom have strong tendency to be oxidized; they are strong reducing agents. • E°ox = -E°red • For any cell, E°cell = E°ox + E°red. • Never multiply half-cell potentials by coefficients used to balance redox reactions.

29. IV. Sample Problem • Calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25 °C. 3Pb2+(aq) + 2Cr(s) 3Pb(s) + 2Cr3+(aq)

30. IV. Predicting Spontaneous Redox Reactions • Positive cell potentials are spontaneous. • We can identify the half-cells in a reaction, look up their potentials, and sum to find overall cell potentials. • Thus spontaneity can be determined for any redox reaction.

31. IV. Sample Problem • Will the following redox reaction be spontaneous under standard conditions? • Zn(s) + Ni2+(aq) Zn2+(aq) + Ni(s) • Zn(s) + Ca2+(aq)  Zn2+(aq) + Ca(s)

32. IV. Sample Problem • Determine whether the following metals dissolve in HCl(aq), HNO3(aq), both, or neither. • Fe(s) • Au(s) • Ag(s)