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Chapter 20 Electrochemistry.

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Presentation on theme: "Chapter 20 Electrochemistry."— Presentation transcript:

1 Chapter 20 Electrochemistry

2 Overview Oxidation-Reduction reactions Voltaic Cells
Balancing Redox Reactions Half-Reaction method Acidic Solution Basic Solution Voltaic Cells Cell EMF--standard reduction potentials Oxidizing & Reducing reagents Spontaneity of Redox reactions

3 Effect of Concentration
Nernst Equation Equilibrium Constants Commercial Voltaic Cells Electrolysis Quantitative Aspects Electrical Work

4 Redox Reactions Involve a transfer of electrons
Oxidation  loss of one or more electron(s) oxidation state will increase Reduction  gain of one or more electron(s) oxidation state will decrease Must occur simultaneously Zn(s) Cu2+(aq)  Zn2+(aq) Cu(s) Zn  Zn2+(aq) e- oxidation oxidation ½ rxn Cu2+(aq) + 2e-  Cu(s) reduction ½ rxn reduction

5 You must know oxidation states: (Review: Section 8.10)
What are the oxidation states of each atom in the following: H2 CO ClO2- HC2H3O2 H 0 C +2, O -2 Cl +3, O -2 H +1, C1 +3, C2 -3, O -2

6 Balancing Redox Reactions
Mass balance must be observed e--transfer must be balanced Simple reactions: Sn Fe3+  Sn Fe2+ Sn2+  Sn e- Fe e-  Fe2+ oxidation ½ rxn x 2 reduction ½ rxn 2Fe e-  2Fe2+ Sn Fe3+  Sn Fe2+

7 Reactions involving H & O in acid:
MnO C2O  Mn CO2 write both ½ reactions MnO  Mn2+ C2O  CO2 mass balance (all except H & O) C2O  CO2 add H2O & H+ to balance O & H 8H MnO  Mn H2O

8 check the balance balance charge by adding electrons
5e H MnO  Mn H2O C2O  CO e- balance electrons transferred 10e H MnO  2Mn H2O 5C2O  CO e- add half reactions 16H+ + 2MnO4-+ 5C2O42-  10CO2 + 2Mn H2O check the balance

9 check balance Reactions in base: MnO4- + CN-  CNO- + MnO2
use exactly the same process CN-  CNO- MnO4-  MnO2 H2O + + 2H+ + 2e- 3e- + 4H+ + + 2H2O since H+ cannot exist in basic solution, add OH- 2OH- + CN-  CNO- + H2O + 2e- 3e- + 2H2O + MnO4-  MnO2 + 4OH- balance electrons transferred & sum 6OH CN-  3CNO- + 3H2O + 6e- 6e- + 4H2O + 2MnO4-  2MnO2 + 8OH- 3CN- + H2O + 2MnO4-  2MnO2 + 3CNO- +2OH- check balance

10 Voltaic Cells A spontaneous redox reaction that does work Anode
electrode at which oxidation occurs loses mass electrons released, sign is negative Cathode electrode at which reduction occurs gains mass electrons consumed, sign is positive

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12 Cell EMF Difference in potential energy of electrons at the anode and cathode Diff. in potential energy per electrical charge measured in volts 1 V = 1 J C Potential difference = EMF, electromotive force Ecell = cell potential = cell voltage Eºcell = cell potential under std. conditions 1 M, 1 atm, 25 ºC

13 Standard reduction potentials
E ºred in tables E ºcell = E ºred (cathode) - E ºred (anode) Based on “standard hydrogen electrode” 2H+(aq, 1M) + 2e-  H2(g, 1atm) E ºred = 0 V Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g) E ºcell = 0.76 V 0.76 V = 0 V - E ºred (anode) Zn2+(aq, 1M) + 2e-  Zn(s) E ºred (anode) = V

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16 Problem: Calculate Eºcell for Anode: 2Al  2Al3+ + 6e-
2Al(s) + 3I2(s)  2Al3+(aq) I-(aq) Anode: 2Al  2Al e- Cathode: 3I e-  6I- Eºcell = E ºred (cathode) - E ºred (anode) E ºcell = V - (-1.66 V) E ºcell = V

17 The more positive the E ºcell the more driving force for the reaction
Note: stoichiometric coefficient does not affect the value of the E ºred (it is an intensive property) E ºox = - E ºred 2Al(s) + 3I2(s)  2Al3+(aq) I-(aq) 2Al  2Al e- E ºox = V 3I e-  6I- E ºred = V E ºcell = E ºox + E ºred = 2.20V The more positive the E ºcell the more driving force for the reaction

18 Oxidizing/Reducing Agents
Oxidizing agents cause oxidation oxidizing agents are reduced the more (+) the E ºred the better the ox. agent Reducing agents cause reduction reducing agents are oxidized the more (-) the E ºred the better the red. agent

19 Which is the better oxidizing agent?
NO H e-  NO + 2H2O E ºred V Ag e-  Ag E ºred V Cr2O H+ + 6e-  2Cr3+ + H2O E ºred V Which is the strongest reducing agent? I e-  2I Eºred V Fe e-  Fe Eºred V MnO4- + 8H+ + 5e-  Mn2+ + 4H2O Eºred V

20 Spontaneity of Redox Reactions
Spontaneous redox rxns have positive potentials Non-spontaneous redox rxns have negative potentials Is this rxn spont. or non-spont.? MnO4- + 8H+ + 5Fe2+  5Fe3+ + Mn2+ + 4H2O Fe2+  Fe3+ + 1e- Eºox = v MnO4- + 8H+ + 5e-  Mn2+ + 4H2O E ºred = v E ºox + E ºred = v Yes

21 EMF & Free Energy If both DG & E are a measure of spontaneity, they must be related DG = - nFE F is Faraday’s constant 1 F = 96,500 J/v mol e- remember: 1 C = 1 J/v n = mol e- transferred In the standard state DGº = - nFEº

22 DG = - (2 mol e-)(-0.083 v)(96,500 J/v mol e-) = + 16 kJ
Calculate the standard free energy change for Hg + 2Fe3+  Hg2+ + 2Fe2+ n = 2 mol electrons transferred Hg  Hg2+ + 2e- Eox = v 2Fe3+ +2e-  2Fe2+ Ered= v Ecell = v DG = - (2 mol e-)( v)(96,500 J/v mol e-) = + 16 kJ

23 Concentration & Cell EMF
Nernst Equation relationship between DG & concentrations DG = DGº + RT ln Q Q = [prod]x/[react]y substitute -nFE for DG E = Eº - (RT/nF) ln Q or E = Eº - (2.303 RT/nF) log Q 2.303 RT/F = v-mol e- at std. temp. E = Eº - (0.0592/n) log Q

24 Calculate the emf that the following cell generates when [Mn2+] = 0
Calculate the emf that the following cell generates when [Mn2+] = 0.10 M & [Al3+] = 1.5 M 2Al + 3Mn2+  2Al3+ + 3Mn Eº = v E = ( v) - ( v/ 6) log [(1.5)2/(0.10)3] E = v when [Mn2+] = 1.5 M & [Al3+] = 0.10 M E = ( v) - ( v/ 6) log [(0.10)2/(1.5)3] E = v

25 Equilibrium Constants
Remember DG = DGº + RT ln Q, if Q = K, then DG = 0, therefore -nFE = 0 and 0 = Eº - (RT/nF) ln K or 0 = Eº - (0.0592/n) log K K can be calculated from cell potentials log K = nE º/0.0592

26 Calculate the equilibrium constant, K, for 2IO3- + 5Cu + 12H+  I2 + 5Cu2+ + 6H2O
Eº = v n = 10 mol e- transferred log K = nEº/0.0592 log K = 145 K = 1 x 10145

27 Voltaic Cells Lead storage battery Dry cell
PbO2 + SO H+ + 2e-  PbSO4 + H2O Pb + SO  PbSO4 + 2e- Ecell = v Dry cell NH4+ + 2MnO2 + 2e-  Mn2O3 + 2NH3 + H2O Zn  Zn2+ + 2e- In an alkaline cell the NH4Cl is replaced with KOH

28 Ni-Cd NiO2 + 2H2O + 2e-  Ni(OH)2 + 2OH- Cd + 2OH-  Cd(OH)2 + 2e- Fuel cells 4e- + O H2O  4OH- 2H2 + 4OH-  H2O

29 Electrolytic Cells Redox reactions that are not spontaneous
Must be driven by an outside source of electrical energy Cathode reduction occurs by sign convention, is negative Anode oxidation occurs by sign convention, is positive

30 Quantitative Aspects Redox reactions occur in stoichiometric relationship to the transfer of electrons Electrons put into a system through electrical energy, can be quantized Coulomb = quantity of charge passing through electrical circuit in 1 s at 1 ampere (A) current Coulomb = (amp) (seconds)

31 Problem: Calculate the mass of Mg formed upon passage of a current of 60.0 A for a period of 4.00 x s. MgCl2  Mg Cl2 Mg e-  Mg 2Cl-  Cl e- we are concerned with the reduction (60.0 A)(4 x 103s)(1C/1 A-s) = 2.4 x 105 C (2.4 x 105 C)(1 mol e-/ 96,500 C) = 2.49 mol e- (2.49 mol e-)(1 mol Mg/2 mol e-) = 1.24 mol Mg (1.24 mol Mg)(24.3 g/mol) = 30.1 Mg

32 Electrical Work DG = wmax DG = - nFE wmax = - nFE
Max work proportional to potential wmax = n F E J = (mol) (C/mol) (J/C) Electrical work = (watt) (time) 1 watt (W) = 1 J/s or watt-s = J 1 kWh = x 106 J


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