CHM 120 CHAPTER 21 Electrochemistry: Chemical Change and Electrical Work

1. CHM 120CHAPTER 21Electrochemistry: Chemical Change and Electrical Work Dr. Floyd Beckford Lyon College

2. REVIEW • Oxidation: the loss of electrons by a species, • leading to an increase in oxidation number of one • or more atoms • Reduction: the gain of electrons by a species, • leading to an decrease in oxidation number of one • or more atoms • Oxidizing agents: the species that is reducedin a • redox reaction

3. Reducing agents: the species that is oxidized in a redox reaction

4. In acidic solution: add H+ or H2O only In basic solution: add OH- or H2O only To balance O To balance H In acidic solution: Add H2O and then Add H+ For each O needed For each H needed In basic solution 1. add 2 OH- to the side needing O and and then 1. add 1 H2O to the side needing H and 2. add 1 H2O to the other side 2. add 1 OH- to the other side

5. THE HALF-REACTION METHOD • This method breaks the overall reaction into its • two components – half-reactions • Each half-reaction is balanced separately and • then added • Use the following guidelines to help • Write as much of the unbalanced net ionic • equation as possible • 2. Decide which atoms are oxidized and which are • reduce – write the two unbalanced half-reactions

6. 3. Balance by inspection all atoms in each half- • reaction except H and O • 4. Use the rules mentioned previously to balance • H and O in each half-reaction • 5. Make equal the number of electrons involved in • both half-reactions • Take a look at the breathalyzer reaction • H+(aq) + Cr2O72-(aq) + C2H5OH(l)  • Cr3+(aq) + C2H4O(l) + H2O(l)

7. Balance the following net ionic equation in basic solution.

9. ELECTROCHEMISTRY • Deals with chemical changes produced by an • electric current and with the production of • electricity by chemical reactions • All electrochemical reactions involve transfer • of electrons and are redox reactions • EChem reactions take place in electrochemical • cell (an apparatus that allows a reaction to • occur through an external conductor)

10. ELECTROCHEMICAL CELLS Two types: 1. Electrolytic cells: - these are cells in which an external electrical source forces a nonspontaneousreaction to occur 2. Voltaic cells: - also called galvanic cells. In these cells spontaneouschemical reactions generate electrical energy and supply it to an external circuit

11. Electric current enters and exits the cell by • electrodes - electrodes are surfaces upon which • oxidation or reduction half-reactions occur • Inert electrodes: - electrodes that don’t react • Two kindsof electrodes: • 1. Cathode: - electrode at whichreduction • occurs (electrons are gained by a species) • 2. Anode: - electrodeat which oxidation occurs • (as electrons are lost by some species)

12. VOLTAIC CELLS • Cells in which spontaneous reactions produces • electrical energy • The two half-cells are separated so that electron • transfer occurs through an external circuit • Each half-cell contains the oxidized and reduced • forms of a species in contact with each other • Half-cells linked by a piece of wire and a salt • bridge

14. A salt bridge has three functions: • 1. It allows electrical contact between the two • half-cells • 2. It prevents mixing of the electrode solutions • 3. It maintains electrical neutrality in each • half-cell as ions flow into and out of the salt • bridge • Point 2 is important – no current would flow if • if both solutions were in the same cell

15. Point 3 is also important – anions flow into the • oxidation half-cell to counter the build-up • of positive charge • Current flow spontaneously from negative to • the positive electrode • In all voltaic cells the anode is negative and the • cathode is positive

17. In voltaic cells, voltage drops as the reaction • proceeds. When voltage = 0, the reaction is at • equilibrium

18. The Silver-Copper cell • Composed of two half-cells: • 1. A strip of copper immersed in 1 M CuSO4 • 2. A strip of silver immersed in 1 M AgNO3 • Experimentally we see: • : - Initial voltage is 0.46 volts • : - The mass of the copper electrode decreases • : - The mass of the silver electrode increases • : - [Cu2+] increases and [Ag+] decreases

19. Cu  Cu2+ + 2e- (oxidation, anode) • 2(Ag+ + e- Ag) (reduction, cathode) • 2Ag+ + Cu  Cu2+ + Ag (Overall cell reaction) • Cu |Cu2+(1.0 M) ||Ag+(1.0 M) | Ag • Notice that in this case the copper electrode is • the anode

20. STANDARD ELECTRODE POTENTIALS • Associated with each voltaic cell is a potential • difference called the cell potential, Ecell • E measures the spontaneity of the cell’s redox • reaction • Higher (more positive) cell potentials indicate a • greater driving force for the reactionas written • All electrode potentials are measured versus the • Standard Hydrogen Electrode (SHE): E° = 0.00 V

21. The E°cell calculated is for the cell operating • under standard state conditions • For electrochemical cell standard conditions • are: • solutes at 1 M concentrations • gases at 1 atm partial pressure • solids and liquids in pure form • All at some specified temperature, usually 298 K

23. The electrode potential for each half-reaction is • written as a reduction process • The more positive the E° value for a half- • reaction the greater the tendency for the reaction • to proceed as written • The more negative the E° value, the more likely • is the reverse of the reaction as written

24. Prediction of Spontaneity 1. First write the HR equation with the more positive (less negative) E° for the reduction along with its potential 2. Write the other HR as anoxidation and include its oxidation potential 3. Balance the electron transfer 4. Add the reduction and oxidation HR and add the corresponding electrode potentials to get the overall cell potential, E°cell

25. Important points to note: • 1. E° for oxidation half-reactions are equal to • but opposite in sign to reduction half-reactions • 2. Half-reaction potentials are the same • regardless of the species’ stoichiometric • coefficient in the balanced equation • E°cell > 0 Forward reaction is spontaneous • E°cell < 0 Backward reaction is spontaneous

26. E°cell, G° and K • From thermodynamics, we know that, • DG° = -RT lnK • We can relate E°cell to free energy for that cell • DG° = -nFE°cell • n = number of moles of e- • So -nFE°cell = -RT lnK and • E°cell = (RT/nF) lnK

27. (Standard state conditions) • Under nonstandard conditions • G = -nFEcell

28. THE NERNST EQUATION • Usually concentrations of reactants differ from • one another and also change during the course • of a reaction • As cell reaction proceeds, cell voltage drops so • that E°cell is different from Ecell • E°cell and Ecell are related by the Nernst • Equation • Ecell = E°cell - (RT/nF) lnQ

29. Ecell = E°cell - (RT/nF) lnQ E = potential under the nonstandard conditions E° = standard potential R = gas constant, 8.314 J/mol.K T = absolute temperature n = number of moles of electrons transferred F = faraday, 96,485 J/V.mol e- Q = reaction quotient

30. BATTERIES • Two type of batteries: • : - Primary batteries cannot be “recharged” • Once all the chemicals are consumed there is • no more chemical reaction • : - Secondary batteries can be regenerated • Most common example is the lead storage • batteryused to power automobiles

31. The Lead Storage Battery • Composed of two alternating groups of Pb • plates; one group contains pure lead (anode) and • the other group contains PbO2 (cathode) • The plates are immersed in 40 % sulfuric acid • During discharge • Pb  Pb2+ + 2e- (oxidation) • Pb2+ + SO42- PbSO4 (precipitation) • Net: Pb + SO42- PbSO4 + 2e- (anode)

32. At the cathode • PbO2 + 4H+ + 2e- Pb2+ + 2H2O (reduction) • Pb2+ + SO42- PbSO4 (precipitation) • Net reaction: • PbO2 + 4H+ + SO42- + 2e- PbSO4 + 2H2O • Adding the HR for the two half-cells, gives • Pb + PbO2 + 4H+ + 2SO42- 2PbSO4 + 2H2O • E°cell = 2.041 V • The battery can be recharged

33. Fuel Cells • These are galvanic cells in which the reactants • are continuously supplied to the cell and the • products are continuously removed • Best known example is the hydrogen-oxygen • fuel cell • Hydrogen is fed into the anode compartment • and oxygen into the cathode compartment

34. Oxygen is reduced at the cathode – porous • carbon doped with metallic catalysts • At the anode hydrogen is oxidized to water • Anode: 2H2(g) + 4OH-(aq)  4H2O(l) + 4e- • Cathode: O2(g) + 2H2O(l) + 4e-  4OH-(aq) • Overall: 2H2(g) + O2(g)  2H2O(g)

35. CORROSION • Ordinary corrosion is a redox process in • which metals are oxidized by oxygen in the • presence of moisture • A point of strain on the surface of the metal • acts as an anode • Areas on the metal surface exposed to air • serves as cathodes

36. Anode: Fe(s)  Fe2+(aq) + 2e- • Cathode: O2(g) + 4H+(aq) + 4e-  2H2O(l) • 4Fe(s) + O2(g) + 4H+(aq)  • 4Fe2+(aq) + 2H2O(l) • 2Fe2+(aq) + 4H2O(l)  Fe2O3•H2O(s) + 6H+ • Rust • Al also undergo corrosion – initial oxidation • is stopped by a layer of Al2O3

37. Corrosion prevention 1. Plating a metal with a thin layer of a less easily oxidized metal 2. Allow a protective film to form naturally on the surface of the metal 3. Galvanizing – coating the metal with zinc 4. Cathodic protection – connecting the metal to a “sacrificial anode”

38. ELECTROLYTIC CELLS • Cells in which an electric current causes a • nonspontaneous reaction to occur – one common • process is called electrolysis • In electrolytic cells the anode is the positive • electrode and the cathode is the negative • electrode • Still : Anode = oxidation; cathode = reduction

39. The Down’ Cell: Electrolysis of molten NaCl • Using graphite inert electrodes the following • observations are made • 1. Chlorine, Cl2, is liberated at one electrode • 2. Sodium metal forms at the other electrode • Explanation • 1. Chlorine is produced at the anode by the • oxidation of Cl- ions

40. 2. Metallic sodium is formed by reducing Na+ ions at the cathode 2Cl- Cl2(g) + 2e- (oxidation, anode HR) 2(Na+ + e- Na(l) (reduction, cathode HR) 2Na+ + 2Cl- 2Na(l) + Cl2(g) Overall cell rxn. ________________________________________ • Electrons used at the cathode are reproduced at • the anode • The reaction is nonspontaneous and electricity • is used to force the reaction to occur

41. Electrolysis of aqueous sodium chloride • In an EChem cell containing aqueous NaCl • : - H2 gas is liberated at one electrode • : - Cl2 gas is liberated at the other electrode • : - Solution at the cathode is basic • Rationalization • : - Chloride ions are oxidized at the anode and • H2O is reduced at the cathode

42. 2Cl- Cl2 + 2e- (oxidation, anode) • 2H2O + 2e- 2OH- + H2 (reduction, cathode) • 2H2O + 2Cl- 2OH- + H2 + Cl2 Overall • Sodium metal is more active than hydrogen • metal and liberates H2 from solution • The hydroxide ions are responsible for the • basicity around the cathode

43. FARADAY’S LAW • States that the amount of substance that • undergoes oxidation or reduction at each • electrode during electrolysis is directly • proportional to the amount of electricity that • passes through the cell • One faraday = the amount of electricity that • reduces or oxidizes 1 equivalent of a substance

44. One equivalent of any substance is the amount • of that substance that supplies or consumes one • mole of electrons • 1F = 1mole of electrons • = 6.022 x 1023 e- • = 96,485 C • C = It • C= charge passed; I = current; • t = time (in seconds)