Topic 12 Electrochemistry

1. Topic 12Electrochemistry

2. Review • Oxidation reduction reactions involve a transfer of electrons. • OIL- RIG • Oxidation Involves Loss • Reduction Involves Gain • LEO-GER • Lose Electrons Oxidation • Gain Electrons Reduction

3. Applications • Moving electrons is electric current. • 8H++MnO4-+ 5Fe+2 +5e-® Mn+2 + 5Fe+3 +4H2O • Helps to break the reactions into half reactions. • 8H++MnO4-+5e-® Mn+2 +4H2O (reduction) • 5(Fe+2® Fe+3 + e-) (oxidation) • In the same mixture it happens without doing useful work, but if separate

4. Connected this way the reaction starts • Stops immediately because charge builds up. H+ MnO4- Fe+2

5. Galvanic Cell - uses a spontaneous redox reaction to produce a current that can be used to do work. Salt Bridge allows current to flow H+ MnO4- Fe+2

6. e- • Electricity travels in a complete circuit • Oxidation occurs at the anode • Reduction occurs at the cathode Anode Cathode H+ MnO4- Fe+2

7. Porous Disk H+ MnO4- Fe+2

8. e- e- e- e- Anode Cathode e- e- Reducing Agent Oxidizing Agent

9. Cell Potential • Oxidizing agent pulls the electrons. • Reducing agent pushes the electrons. • The push or pull (“driving force”) is called the cell potential Ecell • Also called the electromotive force (emf) • Unit is the volt(V) = 1 joule of work/coulomb of charge (measured with a voltmeter) • A coulomb is the SI unit of quantity of electricity, (the charge transferred in one second with a constant current of one ampere)

10. 0.76 e- H2 in Pt metal Zn metal Cathode Anode H+ Cl- Zn+2 SO4-2 1 M ZnSO4 1 M HCl

11. Standard Hydrogen Electrode • This is the reference all other oxidations are compared to • Eº = 0 • º indicates standard states of 25ºC, 1 atm, 1 M solutions. H2 in H+ Cl- 1 M HCl

12. Standard Reduction Potentials • It is universally accepted that the half-reaction potential for 2H+ + 2e- → H2 assigned a value of zero volts. • The overall cell potential (Ecell) is then assigned to the other half reaction. • This reaction is always written as a reduction potential • Table 17.1 has a list of the most common Standard Reduction Potentials we will use.

13. Finding Cell Potentials using Standard Reduction Potentials • Fe3+(aq) + Cu(s)®Cu +2(aq) + Fe 2+(aq) • Two half reactions: • Fe3+ + e- → Fe2+E o = 0.77 V • Cu2+ + 2e- → Cu Eo = 0.34 V • Two rules apply: • The half reaction with the largest potential is written as a reduction, the other must be reversed (change its sign). • The number of electrons lost must equal the number of electrons gained (multiply half reactions). • Do not multiply the standard potentials! They do not change!

14. The total cell potential is the sum of the potential at each electrode. • Eocell = Eo(cathode) +Eo(anode) • Eocell = 0.77 V + (– 0.34 V) = 0.43 V • We can look up reduction potentials in table 17.1.

15. Line Notation • solid½Aqueous½½Aqueous½solid • Anode on the left½½Cathode on the right • Single line different phases. • Double line porous disk or salt bridge. • If all the substances on one side are aqueous, a platinum electrode is indicated. • For the last reaction • Cu(s)½Cu+2(aq)½½Fe+2(aq),Fe+3(aq)½Pt(s)

16. Galvanic Cell • The reaction always runs spontaneously in the direction that produced a positive cell potential. • Four things for a complete description. • Cell Potential • Direction of flow • Designation of anode and cathode • Nature of all the components- electrodes and ions

17. Practice • Completely describe the galvanic cell based on the following half-reactions under standard conditions. • MnO4- + 8 H+ +5e-® Mn+2 + 4H2O • Eº = 1.51 • Fe+2 +2e-® Fe(s) Eº = -0.44V

18. Potential, Work and DG • emf = potential (V) = work (J) / Charge(C) • E= work done by system / charge • E = -w/q • Charge is measured in coulombs. • -w = qE • Faraday = 96,485 C/mol e- • q = nF = moles of e- x charge/mole e- • w = -qE = -nFE= DG

19. Potential, Work and DG • DGº = -nFEº • if Eº < 0, then DGº > 0 nonspontaneous • if Eº > 0, then DGº < 0 spontaneous • In fact, reverse is spontaneous. • Calculate DGº for the following reaction: • Cu+2(aq)+ Fe(s) ® Cu(s)+ Fe+2(aq) • Fe+2(aq)+ e-® Fe(s) Eº = 0.44 V • Cu+2(aq)+2e-® Cu(s) Eº = 0.34 V

20. Cell Potential and Concentration • Qualitatively - Can predict direction of change in E from LeChâtelier. • 2Al(s) + 3Mn+2(aq) ® 2Al+3(aq) + 3Mn(s) • Eocell = 0.48v • Predict if Ecell will be greater or less than Eºcell if [Al+3] = 1.5 M and [Mn+2] = 1.0 M • if [Al+3] = 1.0 M and [Mn+2] = 1.5M • if [Al+3] = 1.5 M and [Mn+2] = 1.5 M

21. The Nernst Equation • DG = DGº +RTln(Q) • -nFE = -nFEº + RTln(Q) • E = Eº - RTln(Q) nF • 2Al(s) + 3Mn+2(aq) ® 2Al+3(aq) + 3Mn(s) Eº = 0.48 V • Consider a cell at 25oC where: • [Mn2+] = 0.50 M and [Al3+] = 1.5 M • Use the Nernst equation on the next slide to solve.

22. The Nernst Equation • E = Eº - 0.0591 log(Q) used at 25oC n • As reactions proceed concentrations of products increase and reactants decrease. • The cell will discharge until it reaches equilibrium. • At this point: Q = K (the equilibrium constant) and Ecell= 0 • At equilibrium, the components in the two cells have the same free energy and ∆G=0. • The Cell no longer has the ability to do work.

23. Example: Describe the cell • VO2+ + 2H+ + e- VO2+ + H2O Eo = 1.00v • Zn2+ + 2e- Zn Eo = -0.76v • T = 25oC • [VO2+] = 2.0M • [H+] = 0.50M • [VO2+] = 1.0 x 10-2 M • [Zn2+] = 1.0 x 10-1 M

24. Batteries are Galvanic Cells • Car batteries are lead storage batteries. • Pb +PbO2 +H2SO4®PbSO4(s) +H2O • Dry Cell Zn + NH4+ +MnO2 ® Zn+2 + NH3 + H2O • Alkaline Zn +MnO2 ® ZnO+ Mn2O3 (in base) • NiCad • NiO2 + Cd + 2H2O ® Cd(OH)2 +Ni(OH)2

25. Corrosion • Rusting - spontaneous oxidation. • Most structural metals have reduction potentials that are less positive than O2 . • Fe ® Fe+2 +2e-Eº= 0.44 V • O2 + 2H2O + 4e- ® 4OH- Eº= 0.40 V • Fe+2 + O2 + H2O ® Fe2 O3 + H+ • Reaction happens in two places.

26. Salt speeds up process by increasing conductivity Water Rust e- Iron Dissolves- Fe ® Fe+2

27. Preventing Corrosion • Coating to keep out air and water. • Galvanizing - Putting on a zinc coat • Has a lower reduction potential, so it is more easily oxidized. • Alloying with metals that form oxide coats. • Cathodic Protection - Attaching large pieces of an active metal like magnesium that get oxidized instead.

28. Electrolysis • Running a galvanic cell backwards. • Put a voltage bigger than the potential and reverse the direction of the redox reaction. • Used for electroplating.

29. 1.10 e- e- Zn Cu 1.0 M Cu+2 1.0 M Zn+2 Cathode Anode

30. A battery >1.10V e- e- Zn Cu 1.0 M Cu+2 1.0 M Zn+2 Cathode Anode

31. Steps: • 1. Current and time → charge • amps (C/s) x time (s) to get coulombs • 2. Quantity of charge → mol e- • coulombs (C) x 1/F (mol e- / C) to get mol e- • 3. Moles of e- → mol of element • mol e- x 1 mol of element / mols of e- (needed to form neutral element from ion) • 4. Moles of element → mass of element • mole of element x molar mass / 1 mol

32. Calculating plating • Have to count charge. • Measure current I (in amperes) • 1 amp = 1 coulomb of charge per second • q = I x t • q/nF = moles of metal • Mass of plated metal

33. What mass of copper is plated out when a current of 10.0 amps is passed for 30.0 minutes through a solution containing Cu2+. • How long must 5.00 amp current be applied to produce 10.5 g of Ag from Ag+

34. Other uses • Electroysis of water. • Separating mixtures of ions. • More positive reduction potential means the reaction proceeds forward. • For metals this is typically gaining electrons and forming the solid metal – this removes the ion from solution. • Ions with the more positive the reduction potential will “plate out” first.