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Redox Reactions and Electrochemistry

Redox Reactions and Electrochemistry. Redox Reactions Oxidation Number Oxidizing and Reducing Reagents Galavanic or Voltaic Cells Anode/Cathode/Salt Bridge Cell Notations Determining Cell Potential/Cell Voltage/Electromotive force (emf) Relating Cell Potential to K and D G 0

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Redox Reactions and Electrochemistry

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  1. Redox Reactions and Electrochemistry • Redox Reactions • Oxidation Number • Oxidizing and Reducing Reagents • Galavanic or Voltaic Cells • Anode/Cathode/Salt Bridge • Cell Notations • Determining Cell Potential/Cell Voltage/Electromotive force (emf) • Relating Cell Potential to K and DG0 • Effect of Concentration on Cell Potential

  2. Redox Reactions and Electrochemistry • Corrosion • Batteries • Fuel Cells • Electrolytic Cells • Calculating amounts of substances reduced or oxidized

  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

  4. 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

  5. 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

  6. 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

  7. 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

  8. 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

  9. Determination of Oxidizing and Reducing Agents • Determine oxidation # for all atoms in both the reactants and products. • Look at same atom in reactants and products and see if oxidation # increased or decreased. • If oxidation # decreased; substance reduced • If oxidation # increased; substance oxidized

  10. Determination of Oxidizing and Reducing Agents, continued • Oxidizing Agent: Substance that oxidizes the other substance by accepting electrons. It is reduced in reaction. • Reducing Agent: Substance that reduces the other substance by donating electrons. It is oxidized in reaction.

  11. Spontaneous Redox ReactionZn(s) + Cu+2 (aq) -> Cu(s) + Zn+2(aq) Zn Cu time Zn+2 Cu+2

  12. <- Gets Larger Gets Smaller ->

  13. 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

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

  15. 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

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

  17. 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

  18. 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

  19. Determining if Redox Reaction is Spontaneous • + E°CELL ; spontaneous reaction • E°CELL = 0; equilibrium • - E°CELL; nonspontaneous reaction More positive E°CELL ; stronger oxidizing agent or more likely to be reduced

  20. 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

  21. Relating E0Cell to DG0 Units work, Joule charge, Coulomb Ecell; Volts Faraday, F; charge on 1 mole e- F = 96485 C/mole charge = nF work = (charge)Ecell = -nFEcell DG = work (maximum) DG = -nFEcell

  22. Relating EoCELL to the Equilibrium Constant, K DG0 = -RT ln K -RT ln K = -nFE0cell DG0 = -nFE0cell

  23. Effect of Concentration on Cell Potential DG =DG0 + RTlnQ -nFEcell= -nFE0cell +RTln Q DG0 = -nFE0cell Ecell= E0cell - RTln Q nF Ecell= E0cell - 0.0257ln Q n Ecell= E0cell – 0.0592log Q n

  24. Corrosion – Deterioration of Metals by Electrochemical Process

  25. Corrosion – Deterioration of Metals by Electrochemical Process

  26. Corrosion – Deterioration of Metals by Electrochemical Process

  27. Cathodic Protection

  28. Abbreviated Standard Reduction Potential Table

  29. 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

  30. 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

  31. 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

  32. 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

  33. 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

  34. 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.

  35. Charge =(Current)(Time) Molar Mass Redox Eqn Faraday’s Constant

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