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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 DG 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 • Corrosion • Batteries • Fuel Cells • Electrolytic Cells • Calculating amounts of substances reduced or oxidized

  2. REDOX REACTION

  3. DEFINITIONS • Oxidation: Loss of electrons. • Reduction: Gain of electrons. LEO says GER • Oxidation: Gain of oxygen • Reduction: Lost of oxygen • Oxidation: Increasing of oxidation number • Reduction: Reducing of oxidation number

  4. CO  CO2(CO oxidized) why? CH3COOH  CH3CHO (CH3COOH reduced) why? H2SO3  H2SO4 ?? HNO3 HNO2 ?? Ca2+ Ca (Ca2+ reduced) why? Na  Na+(Na oxidized) why? Fe3+  Fe2+ Mn2+ MnO4–(Mn2+ oxidized) why? Cl2 2 Cl–(Cl2 reduced) why? H2O2 H2O ?? NaH H2 ??

  5. LEO says GER : Lose Electrons = Oxidation Sodium is oxidized Na0 Na+1 + 1e- Gain Electrons = Reduction Chlorine is reduced Cl0 + 1e– Cl–

  6. Rules for Assignment of Oxidation Number (ON) 1) The ON of all pure elements is zero. 2) The ON of His +1, except in hydrides, where it is -1. 3) The ON of O is -2, except in peroxides, where it is -1. 4) The algebraic sum of ON must equal zero for a neutral molecule or the charge on an ion.

  7. Variable Oxidation Number of Elements • Sulfur:SO42-(+6), SO32-(+4), S(0), FeS2(-1), H2S(-2) • Carbon:CO2(+4), C(0), CH4(-4) • Nitrogen:NO3-(+5), NO2-(+3), NO(+2), N2O(+1), N2(0), NH3(-3) • Iron:Fe2O3(+3), FeO(+2), Fe(0) • Manganese:MnO4-(+7), MnO2(+4), Mn2O3(+3), MnO(+2), Mn(0) • Copper:CuO(+2), Cu2O(+1), Cu(0) • Tin:SnO2(+4), Sn2+(+2), Sn(0) • Uranium:UO22+(+6), UO2(+4), U(0) • Arsenic:H3AsO40(+5), H3AsO30(+3), As(0), AsH3(-1) • Chromium:CrO42-(+6), Cr2O3(+3), Cr(0) • Gold:AuCl4-(+3), Au(CN)2-(+1), Au(0)

  8. BALANCING OVERALL REDOX REACTIONS Example balance the redox reaction below: Fe + Cl2 Fe3+ + Cl- Step 1: Assign oxidation number, Fe0 + Cl20 Fe3+ + Cl- Step 2: Determine number of electrons lost or gained by reactants. Fe0 + Cl20 Fe3+ + Cl-  3e- 2e- Step 3: Cross multiply. 2Fe + 3Cl20 2Fe3+ + 6Cl-

  9. may be written as the sum of two half-cell reactions: 2Fe  2Fe3+ + 6e- (oxidation) 3Cl20+ 6e- 6Cl- (reduction) All overall redox reactions can be expressed as the sum of two half-cell reactions, one a reduction and one an oxidation. The overall reaction: 2Fe + 3Cl20 2Fe3+ + 6Cl-

  10. Another example - balance the redox reaction: FeS2 + O2 Fe(OH)3 + SO42- Fe+2S20+ O20 Fe+3(OH)3 + S+6O42-   15e- 4e- 4FeS2 + 15O2 Fe(OH)3 + SO42- 4FeS2 + 15O2 4Fe(OH)3 + 8SO42- 4FeS2 + 15O2 +14H2O  4Fe(OH)3 + 8SO42- + 16H+ This reaction is the main cause of acid generation in drainage from sulfide ore deposits. Note that we get 4 moles of H+ for every mole of pyrite oxidized!

  11. 2Mg 2Mg2+ + 4e- Electrochemistry: Interconversion of electrical and chemical energy using redox reactions • Oxidation Half-Reaction:Oxidation Involves Loss • of electrons • Reduction Half-Reaction:Reduction Involves Gain of electrons O2 + 4e- 2O2- Net Redox Reaction: 2Mg + O2 2 Mg+2 + 2 O-2

  12. 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 formonoatomic 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

  13. 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.

  14. 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

  15. 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- IF7 NaIO3 O = -2 F = -1 Na = +1 H = +1 7x(-1) + ? = 0 O = -2 I = +7 3x(-2) + 1 + ? = 0 3x(-2) + 1 + C = -1 I = +5 C = +4

  16. Determination of Oxidizing and Reducing Agents Determine oxidation number for all atoms in both the reactants and products. Look at same atom in reactants and products and see if oxidation number increased or decreased. If oxidation number decreased:substance reducedOxidizing Agent If oxidation number increased; substance oxidized Reducing Agent

  17. Zn(s) + Cu+2 (aq) -> Cu(s) + Zn+2(aq) Spontaneous Redox Reaction Zn Cu time Zn+2 Cu+2

  18. Gets Larger Gets Smaller

  19. 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 nonspontaneousredox reaction.

  20. Voltaic Cell 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

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

  22. 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

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

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

  25. 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)

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

  27. 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

  28. Relating E0Cell to G0 Units Work: Joule Charge (Q): Coulomb Ecell : Volts Charge = nF Faraday (F): charge on 1 mole e- F = 96485 C/mole Work = (charge)Ecell = -nFEcell G = work (maximum) G = -nFEcell

  29. Relating EoCELL to the Equilibrium Constant, K G0 = -RT ln K -RT ln K = -nFE0cell G0 = -nFE0cell

  30. Effect of Concentration on Cell Potential G = G0 + RTlnQ -nFEcell = -nFE0cell +RTln Q G0 = -nFE0cell Ecell = E0cell - RTln Q nF Ecell= E0cell - 0.0257ln Q n Ecell= E0cell – 0.0592log Q n

  31. Corrosion – Deterioration of Metals by Electrochemical Process

  32. Cathodic Protection

  33. Batteries Dry cell A: Zn (s) Zn2+ (aq) + 2e- C: 2NH4+ (aq) + 2MnO2 (s) + 2e- Mn2O3 (s) + 2NH3 (aq) + H2O (l) Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)

  34. 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)

  35. 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)

  36. 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

  37. 2H2 (g) + 4OH- (aq) 4H2O (l) + 4e- 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 Anode: Cathode:

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