Redox Reactions and Electrochemistry

1. Redox Reactions and Electrochemistry 9.3 Oxidation Numbers ( only first half, assigning ox numbers) 10.1 Galavanic or Voltaic Cells • Anode/Cathode/Salt Bridge • Cell Notations • Determining Cell Potential/Cell Voltage/Electromotive force (emf)

2. Application • Corrosion • Batteries • Fuel Cells • Electrolytic Cells

3. 2Mg 2Mg2+ + 4e- O2 + 4e- 2O2- Electrochemistry: Interconversion of electrical and chemical energy using redox reactions Redox (Oxidation-Reduction) Reaction: Type of electron transfer reaction. One substance gives up electrons; the other accepts electrons. OIL RIG • Oxidation Half-Reaction;Oxidation Involves Loss of electrons • Reduction Half-Reaction;Reduction Involves Gain of electrons Net Redox Rxn; 2Mg + O2 -> 2 Mg+2 + 2 O-2

5. Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred to the more electronegative atom. • Oxidation number equals ionic charge for monoatomic ions in ionic compound CaBr2; Ca = +2, Br = -1 2. Metal ions in Family A have one, positive oxidation number; Group IA metals are +1, IIA metals are +2 Li+, Li = +1; Mg+2, Mg = +2 4.4

6. Oxidation number,continued The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred to the more electronegative atom. 3. The oxidation number of a transition metal ion is positive, but can vary in magnitude. • Nonmetals can have a variety of oxidation numbers,both positive and negative numbers which can vary in magnitude. 5. Free elements (uncombined state) have an oxidation number of zero. Each atom in O2, F2, H2, Cl2, K, Be has the same oxidation number; zero. 4.4

7. 6. The oxidation number of fluorine is always–1. (unless fluorine is in elemental form, F2) • 7. 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. IF; F= -1; I = +1 8. The oxidation number of hydrogen is +1except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1or when it’s in elemental form (H2; oxidation # =0). HF; F= -1, H= +1 NaH; Na= +1, H = -1

8. Oxidation numbers of all the atoms in HCO3- ? 9. The oxidation number of oxygen is usually–2. In H2O2 and O22- it is –1, in elemental form (O2 or O3) it is 0. H2O ; H=+1, O= -2 SO3; O = -2; S = +6 HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4 4.4

9. Oxidation numbers of all the elements in the following ? IF7 F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 O = -2 K = +1 O = -2 Na = +1 3x(-2) + 1 + ? = 0 7x(-2) + 2x(+1) + 2x(?) = 0 I = +5 Cr = +6 4.4

10. +2 +5 -8 +1 +5 -6 +2 +12 -14 -4 +6 -2 +1 -1 +1 +5 -2 +1 +5 -2 +1 +6 -2 -2 +1 -2 +1 -1 H2PO4– HNO3 K2Cr2O7 C2H6O AgI

11. Application to Electrochemistry • An electric cell converts chemical energy into electrical energy • Alessandro Volta invented the first electric cell but got his inspiration from Luigi Galvani. Galvani’s crucial observation was that two different metals could make the muscles of a frog’s legs twitch. Unfortunately, Galvani thought this was due to some mysterious “animal electricity”. It was Volta who recognized this experiment’s potential. • An electric cell produces very little electricity, so Volta came up with a better design: • A battery is defined as two or more electric cells connected in series to produce a steady flow of current • Volta’s first battery consisted of several bowls of brine (NaCl(aq)) connected by metals that dipped from one bowl to another • His revised design, consisted of a sandwich of two metals separated by paper soaked in salt water.

12. Alessandro Volta’s invention was an immediate technological success because it produced electric current more simply and reliably than methods that depended on static electricity. It also produced a steady electric current –something no other device could do.

13. Electric cells are composed of two electrodes – solid electrical conductors and at least one electrolyte (aqueous electrical conductor) • In current cells, the electrolyte is often a moist paste (just enough water is added so that the ions can move). Sometimes one electrode is the cell container. • The positive electrode is defined as the cathode and the negative electrode is defined as the anode • The electrons flow through the external circuit from the anode to the cathode.To test the voltage of a battery, the red(+) lead is connected to the cathode (+ electrode), and the black(-) lead is connected to the anode (- electrode)

14. Electric potential difference is like the potential energy difference between 1kg of water at the top of the dam and 1kg of water at the bottom of the dam. Electric current (or flow of electrons) is like the flow of the water. A larger drain would be like a larger electric cell, allowing more water (or electrons) to flow. • A voltmeter is a device that measures the energy difference, per unit charge, between any two points in an electric circuit (called electric potential difference) • I.e. A 9V battery releases 6X as much energy compared with the electrons from a 1.5V battery. • The voltage of a cell depends mainly on the chemical composition of the reactants in the cell • An ammeter is a device that measures the rate of flow of charge past a point in an electrical circuit (called electric current) • The larger the electric cell, the greater current that can be produced

15. Secondary cells can be recharged using electricity, but are expensive Primary cells cannot be recharged, but are relatively inexpensive Fuel cells produce electricity by the reaction of a fuel that is continuously supplied. More efficient, and used for NASA vehicles, but still too expensive for general or commercial applications

17. <- Gets Larger Gets Smaller ->

18. Voltaic Cell Animation AnOx or both vowels Red Cat or both consonants Anode; Site of Oxidation Cathode; Site of Reduction Direction of electron flow; anode to cathode (alphabetical) Salt Bridge; Maintains electrical neutrality + ion migrates to cathode - ion migrates to anode

19. Cell Notation • Anode • Salt Bridge • Cathode Anode | Salt Bridge | Cathode | : symbol is used whenever there is a different phase

20. Zn (s) + Cu2+(aq) Cu (s) + Zn2+(aq) Cell Notation [Cu2+] = 1 M & [Zn2+] = 1 M Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s) cathode anode More detail.. Zn (s)| Zn+2 (aq, 1M)| K(NO3) (saturated)|Cu+2(aq, 1M)|Cu(s) Salt bridge anode cathode 19.2

21. Zn (s) + 2 H+(aq) -> H2 (g) + Zn+2 (aq) K(NO3) Zn(s)| Zn+2|KNO3|H+(aq)|H2(g)|Pt

22. Electrochemical Cells • The difference in electrical potential between the anode and cathode is called: • cell voltage • electromotive force (emf) • cell potential UNITS: Volts Volt (V) = Joule (J) Coulomb, C 19.2

27. 2e- + 2H+ (1 M) H2 (1 atm) Standard Electrode Potentials Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm. Reduction Reaction E0= 0 V Standard hydrogen electrode (SHE) 19.3

29. Determining if Redox Reaction is Spontaneous • + E°CELL ; spontaneous reaction • - E°CELL; nonspontaneous reaction More positive E°CELL ; stronger oxidizing agent or more likely to be reduced

30. E0 is for the reaction as written • The half-cell reactions are reversible • The sign of E0changes when the reaction is reversed • Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0 • The more positive E0 the greater the tendency for the substance to be reduced 19.3

31. Standard Potentials (cont.) • If in constructing an electrochemical cell, you need to write the reaction as a oxidation instead of a reduction, the sign of the 1/2 cell potential changes. Zn+2 + 2e- Zn E° = -0.76 V Zn Zn+2 + 2e- E° = +0.76 V • 1/2 cell potentials are intensive variables. As such, you do NOT multiply them by any coefficients when balancing reactions.

32. Writing Galvanic Cells (cont.) Shorthand Notation Zn|Zn+2||Cu+2|Cu Anode Cathode Salt bridge

33. Predicting Galvanic Cells • Given two 1/2 cell reactions, how can one construct a galvanic cell? • Need to compare the reduction potentials of the two half cells. • Turn the reaction for the weaker reduction (smaller E°1/2) and turn it into an oxidation. This reaction will be the anode, the other the cathode.

34. Predicting Galvanic Cells (cont.) • Example. Describe a galvanic cell based on the following: Ag+ + e- Ag E°1/2 = 0.80 V Fe+3 + e- Fe+2 E°1/2 = 0.77 V Weaker reducing agent – turn it around Ag+ + Fe+2 Ag + Fe+3 E°cell = 0.03 V E°cell > 0….cell is galvanic

35. Another Example • For the following reaction, identify the two half cells, and use these half cells to construct a galvanic cell 3Fe+2(aq) Fe(s) + 2Fe+3(aq) oxidation +2 0 +3 reduction Fe+2(aq) + 2e- Fe(s) E° = -0.44 V Fe+3(aq) + e- Fe+2(aq) E° = +0.77 V

36. Another Example (cont.) weaker reduction – turn it around Fe+2(aq) + 2e- Fe(s) E° = -0.44 V Fe+3(aq) + e- Fe+2(aq) E° = +0.77 V 2 x 2Fe+3(aq) + Fe(s) 3Fe+2(aq) E°cell = 1.21 V E° = +0.44 V Fe(s) Fe+2(aq) + 2e-

37. Corrosion – Deterioration of Metals by Electrochemical Process

38. Corrosion – Deterioration of Metals by Electrochemical Process

40. Cathodic Protection

41. Abbreviated Standard Reduction Potential Table

43. Zn (s) Zn2+ (aq) + 2e- Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s) Batteries Dry cell Leclanché cell Anode: 2NH4+(aq) + 2MnO2(s) + 2e- Mn2O3(s) + 2NH3(aq) + H2O (l) Cathode: 19.6

44. Batteries Mercury Battery Anode: Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e- Cathode: HgO (s) + H2O (l) + 2e- Hg (l) + 2OH-(aq) Zn(Hg) + HgO (s) ZnO (s) + Hg (l) 19.6

45. PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- PbSO4(s) + 2H2O (l) Batteries Lead storage battery Anode: Pb (s) + SO42- (aq) PbSO4 (s) + 2e- Cathode: Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq) 2PbSO4 (s) + 2H2O (l) 19.6

46. Fuel Cell vs. Battery • Battery; Energy storage device • Reactant chemicals already in device • Once Chemicals used up; discard (unless rechargeable) • Fuel Cell; Energy conversion device • Won’t work unless reactants supplied • Reactants continuously supplied; products continuously removed

47. 2H2 (g) + 4OH- (aq) 4H2O (l) + 4e- Anode: Cathode: O2(g) + 2H2O (l) + 4e- 4OH-(aq) 2H2 (g) + O2 (g) 2H2O (l) Fuel Cell A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning

48. Types of Electrochemical Cells • Voltaic/Galvanic Cell; Energy released from spontaneous redox reaction can be transformed into electrical energy. • Electrolytic Cell; Electrical energy is used to drive a nonspontaneous redox reaction.