1. Electrochemistry Chapter 4 4.4 and 4.8 Chapter 19 19.1-19.5 and 19.8

2. Oxidation-Reduction Reactions Electron Transfer Reactions Many redox reactions take place in water. Half-reactions OILRIG

3. Oxidation-Reduction Reactions 2Mg (s) + O 2 (g) 2MgO (s) 2Mg 2Mg 2+ + 4e - Oxidation half-reaction (lose e - ) O 2 + 4e - 2O 2- Reduction half-reaction (gain e - ) 2Mg + O 2 + 4e - 2Mg 2+ + 2O 2- + 4e - 2Mg + O 2 2MgO

4. Oxidation-Reduction Reactions Oxidizing Agent- a substance that accepts electrons from another substance, causing the other substance to be oxidized. Reducing Agent- a substance that donates electrons to another substance, causing it to become reduced.

5. Oxidation-Reduction Reactions Zn (s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu (s)

6. Oxidation-Reduction Reactions Zn (s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu (s) Zn Zn 2+ + 2e - Zn is oxidized Zn is the reducing agent Cu 2+ + 2e - CuCu 2+ is reduced Cu 2+ is the oxidizing agent

7. Oxidation-Reduction Reactions Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu (s) + 2AgNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2Ag (s) Cu Cu 2+ + 2e - Ag + + 1e - AgAg + is reduced Ag + is the oxidizing agent

8. Oxidation Number Oxidation Number- the number of charges the atom would have in a molecule (or an ionic compound) if electrons were transferred completely. H 2 (g) + Cl 2 (g) → 2HCl (g) 00 +1 -1

9. Assigning Oxidation Numbers 1.Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H 2, O 2, P 4 = 0 2.In monatomic ions, the oxidation number is equal to the charge on the ion. Li +, Li = +1; Fe 3+, Fe = +3; O 2-, O = -2 3.The oxidation number of oxygen is usually –2. In H 2 O 2 and O 2 2- it is –1.

10. Assigning Oxidation Numbers 4.The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. 5.Group IA metals are +1, IIA metals are +2 and fluorine is always –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.

11. Assigning Oxidation Numbers Oxidation numbers of all the elements in HCO 3 - ? HCO 3 - O = -2H = +1 3x(-2) + 1 + ? = -1 C = +4

13. Assigning Oxidation Numbers

14. Redox Titrations Oxidizing agent/Reducing Agent Equivalence point reached when reducing agent is completely oxidized by the oxidizing agent. Indicator is needed

15. Redox Titrations

18. Electrochemistry Branch of Chemistry that deals with interconversion of electrical and chemical energy. Importance  Use of electricity in everyday life  Interconversion of energy  Study of electrochemical processes

19. Electrochemistry Electrochemical processes are oxidation-reduction reactions in which: the energy released by a spontaneous reaction is converted to electricity or electrical energy is used to cause a nonspontaneous reaction to occur

20. Balancing Redox Equations Simple reactions easy to balance Complex reactions  Usually encountered in laboratory  Can include CrO 4 2-, Cr 2 O 7 2-, MnO 4 -, NO 3 - and SO 4 2-  Follow balancing guidelines

21. Balancing Redox Reactions Guidelines  Step 1: Write the unbalanced equation for the reaction in ionic form.  Step 2: Separate the equation into two half-reactions  Step 3: Balance each half-reaction for number and type of atoms and charges. (Look at medium)  Step 4: Add the two half-reactions together and balance the final equation by inspection. The electrons on both sides must cancel. (Be sure they are equal)  Step 5: Verify that the equation contains the same type and numbers of atoms and the same charges on both sides of the equation.

22. Balancing Redox Equations 1.Write the unbalanced equation for the reaction ion ionic form. The oxidation of Fe 2+ to Fe 3+ by Cr 2 O 7 2- in acid solution? Fe 2+ + Cr 2 O 7 2- Fe 3+ + Cr 3+ 2.Separate the equation into two half-reactions. Oxidation: Cr 2 O 7 2- Cr 3+ +6+3 Reduction: Fe 2+ Fe 3+ +2+3 3.Balance the atoms other than O and H in each half-reaction. Cr 2 O 7 2- 2Cr 3+

23. Balancing Redox Reactions 4.For reactions in acid, add H 2 O to balance O atoms and H + to balance H atoms. Cr 2 O 7 2- 2Cr 3+ + 7H 2 O 14H + + Cr 2 O 7 2- 2Cr 3+ + 7H 2 O 5.Add electrons to one side of each half-reaction to balance the charges on the half-reaction. Fe 2+ Fe 3+ + 1e - 6e - + 14H + + Cr 2 O 7 2- 2Cr 3+ + 7H 2 O 6.If necessary, equalize the number of electrons in the two half- reactions by multiplying the half-reactions by appropriate coefficients. 6Fe 2+ 6Fe 3+ + 6e - 6e - + 14H + + Cr 2 O 7 2- 2Cr 3+ + 7H 2 O

24. Balancing Redox Reactions 7.Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel. 6e - + 14H + + Cr 2 O 7 2- 2Cr 3+ + 7H 2 O 6Fe 2+ 6Fe 3+ + 6e - Oxidation: Reduction: 14H + + Cr 2 O 7 2- + 6Fe 2+ 6Fe 3+ + 2Cr 3+ + 7H 2 O 8.Verify that the number of atoms and the charges are balanced. 14x1 – 2 + 6x2 = 24 = 6x3 + 2x3 9.For reactions in basic solutions, add OH - to both sides of the equation for every H + that appears in the final equation.

27. Galvanic Cells Oxidation-Reduction Reaction Oxidizing agent and Reducing agent are not in contact External conducting medium Galvanic Cell/Volatic Cell- the experimental apparatus for generating electricity through the use of a spontaneous reaction.

28. Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. e – e – Zinc anode Copper cathode Salt bridge Cotton plugs ZnSO 4 solution Zn is oxidized to Zn 2+ at anode. 2e – + Cu 2+ (aq) Cu(s) CuSO 4 solution Zn 2+ SO 4 2– Cl – K + Voltmeter (aq) + Cu(s) Zn 2+ 2e – Cu 2+ Cu SO 4 2– Zn(s) Zn 2+ (aq) + 2 e – Net reaction Cu Cu 2+ is reduced to Cu at cathode. Zn(s) + Cu 2+ (aq) Zn 2+ 2+

29. Galvanic Cells Cell Voltage- the difference in electrical potential between the anode and the cathode. Measured by a voltmeter. Electromotive force (emf) and Cell Potential Voltage of a cell depends on:  Composition of electrode and ions  Concentration of ions  Temperature

30. Galvanic Cells Cell Diagram Zn (s) + Cu 2+ (aq) Cu (s) + Zn 2+ (aq) Zn (s) | Zn 2+ (1 M) || Cu 2+ (1 M) | Cu (s) anodecathode

31. Standard Reduction Potentials Used to calculate the emf of a galvanic cell. Use the values from anode and cathode Standard reduction potential (E 0 ) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm. Standard emf (E 0 ) cell E 0 = E cathode - E anode cell 00

32. Standard Reduction Potentials Zn (s) | Zn 2+ (1 M) || H + (1 M) | H 2 (1 atm) | Pt (s) Anode (oxidation): Zn (s) Zn 2+ (1 M) + 2e - Cathode (reduction): 2e - + 2H + (1 M) H 2 (1 atm) Zn (s) + 2H + (1 M) Zn 2+ + H 2 (1 atm)

33. Standard Reduction Potentials E 0 = 0.76 V cell Standard emf (E 0 ) cell 0.76 V = 0 - E Zn /Zn 0 2+ E Zn /Zn = -0.76 V 0 2+ Zn 2+ (1 M) + 2e - Zn E 0 = -0.76 V E 0 = E H /H - E Zn /Zn cell 00 + 2+ 2 E 0 = E cathode - E anode cell 00 Zn (s) | Zn 2+ (1 M) || H + (1 M) | H 2 (1 atm) | Pt (s)

34. Standard Reduction Potentials 19.3 Pt (s) | H 2 (1 atm) | H + (1 M) || Cu 2+ (1 M) | Cu (s) 2e - + Cu 2+ (1 M) Cu (s) H 2 (1 atm) 2H + (1 M) + 2e - Anode (oxidation): Cathode (reduction): H 2 (1 atm) + Cu 2+ (1 M) Cu (s) + 2H + (1 M) E 0 = E cathode - E anode cell 00 E 0 = 0.34 V cell E cell = E Cu /Cu – E H /H 2++ 2 000 0.34 = E Cu /Cu - 0 0 2+ E Cu /Cu = 0.34 V 2+ 0

37. Tips E 0 is for the reaction as written The more positive E 0 the greater the tendency for the substance to be reduced The half-cell reactions are reversible The sign of E 0 changes when the reaction is reversed Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E 0

39. Spontaneity of Redox Reactions Relationship between Eº cell, ΔGº and K.  G 0 = -nFE cell 0 n = number of moles of electrons in reaction F = 96,500 J V mol = 96,500 C/mol = 0.0257 V n ln K E cell 0 = 0.0592 V n log K E cell 0

40. Spontaneity of Redox Reactions

44. Relationships between Eº cell, ΔGº and K

45. The Effect of Concentration on Cell Emf Not all reactions in galvanic cells can occur under standard state conditions. Nernst Equation- equation used to calculate cell potential under nonstandard state conditions. - 0.0257 V n ln Q E 0 E = - 0.0592 V n log Q E 0 E =

46. Will the following reaction occur spontaneously at 25 0 C if [Fe 2+ ] = 0.60 M and [Cd 2+ ] = 0.010 M? Fe 2+ (aq) + Cd (s) Fe (s) + Cd 2+ (aq) 2e - + Fe 2+ 2Fe Cd Cd 2+ + 2e - Oxidation: Reduction: n = 2 E 0 = -0.44 – (-0.40) E 0 = -0.04 V E 0 = E Fe /Fe – E Cd /Cd 00 2+ - 0.0257 V n ln Q E 0 E = - 0.0257 V 2 ln -0.04 VE = 0.010 0.60 E = 0.013 E > 0Spontaneous 19.5

49. Concentration Cells Galvanic cells consisting of the same type of electrodes in the same solution. Each solution is a different concentration. Zn (s) | Zn 2+ (0.10M) || Zn 2+ (1.0M) | Zn (s) E = Eº - 0.0257V/ 2 x ln [Zn 2+ ] dil / [Zn 2+ ] conc E = 0- 0.0257V/2 ln (0.10/1.0) E = 0.0296 V

50. Concentration Cells Importance  Unequal ion concentrations  Membrane potential  Propagation

51. Electrolysis Electrolysis- is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur. Electrolytic Cell- an apparatus for carrying out electrolysis.