Electrochemistry Chapter 19
Electron Transfer Reactions <ul><li>Electron transfer reactions are  oxidation-reduction  or  redox  reactions. </li></ul>...
Oxidation  half-reaction (lose e - ) Reduction  half-reaction (gain e - ) 19.1 <ul><li>Electrochemical  processes are oxid...
Terminology for Redox Reactions <ul><li>OXIDATION —loss of electron(s) by a species; increase in oxidation number; increas...
You can’t have one… without the other! <ul><li>Reduction (gaining electrons) can’t happen without an oxidation to provide ...
Another way to remember <ul><li>OIL RIG </li></ul>xidation s ose eduction s ain
Review of Oxidation numbers The charge the atom would have in a molecule (or an ionic compound) if electrons were complete...
<ul><li>The oxidation number of hydrogen is  +1   except  when it is bonded to metals in binary compounds.  In these cases...
Balancing Redox Equations 19.1 <ul><li>Write the unbalanced equation for the reaction ion ionic form. </li></ul>The oxidat...
Balancing Redox Equations <ul><li>For reactions in acid, add H 2 O to balance O atoms and H +  to balance H atoms. </li></...
Balancing Redox Equations <ul><li>Add the two half-reactions together and balance the final equation by inspection.  The n...
<ul><li>To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an...
Galvanic Cells 19.2 spontaneous redox reaction anode oxidation cat hode red uction - +
Galvanic Cells 19.2 <ul><li>The difference in electrical potential between the anode and cathode is called: </li></ul><ul>...
Standard Electrode Potentials 19.3 Zn ( s )  |  Zn 2+  (1  M )  || H +  (1  M ) | H 2  (1 atm) | Pt ( s ) Anode (oxidation...
Standard Electrode Potentials 19.3 Standard reduction potential (E 0 )  is the voltage associated with a  reduction reacti...
19.3 <ul><li>E 0  is for the reaction as written </li></ul><ul><li>The more positive  E 0  the greater the tendency for th...
19.3 2 Standard Electrode Potentials Zn ( s )  |  Zn 2+  (1  M )  || H +  (1  M ) | H 2  (1 atm) | Pt ( s ) If the reactio...
Standard Electrode Potentials 19.3 Pt ( s )  |  H 2  (1  atm )  | H +  (1  M ) || Cu 2+  (1  M ) | Cu ( s ) Anode (oxidati...
What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0  M  Cd(NO 3 ) 2  solution and a Cr ele...
19.4 Spontaneity of Redox Reactions  G = - nFE cell n  = number of moles of electrons in reaction = 96,500 C/mol  G 0  =...
Spontaneity of Redox Reactions 19.4
Oxidation: Reduction: 19.4 n  = 2 K  = 1.23 x 10 -42 0 0 2+ + 2e -  + Fe 2+  Fe 2Ag  2Ag +  + 2e - What is the equilibrium...
The Effect of Concentration on Cell Emf  G =   G 0  +  RT  ln  Q  G = - nFE - nFE  = - nFE 0   +  RT  ln  Q Nernst equa...
Oxidation: Reduction: n  = 2 E 0  =  E Fe  /Fe  +  E Cd /Cd 0 0 2+ 2+ E  = 0.013 E  > 0 Spontaneous 19.5 Will the followin...
Charging a Battery When you charge a battery, you are forcing the electrons backwards (from the + to the -).  To do this, ...
Batteries 19.6 Leclanch é cell Dry cell Anode:  Cathode: Zn (s)  Zn 2+  ( aq ) + 2e - 2NH 4   ( aq )  + 2MnO 2   ( s )  + ...
Batteries Anode:   Cathode: Mercury Battery 19.6 Zn(Hg) + 2OH -  ( aq )  ZnO ( s ) + H 2 O ( l ) + 2e - HgO  ( s )  + H 2 ...
Batteries 19.6 Anode:   Cathode: Lead storage battery PbO 2   ( s )  + 4H +   ( aq )  + SO 2-   (aq )  +  2e -   PbSO 4   ...
Batteries 19.6 Solid State Lithium Battery
Batteries 19.6 A  fuel cell  is an electrochemical cell that requires a continuous supply of reactants to keep functioning...
Corrosion 19.7
Cathodic Protection of an Iron Storage Tank 19.7
19.8 Electrolysis  is the process in which electrical energy is used to cause a  nonspontaneous  chemical reaction to occur.
Electrolysis of Water 19.8
Chemistry In Action:  Dental Filling Discomfort Hg 2  /Ag 2 Hg 3   0.85 V 2+ Sn  /Ag 3 Sn  -0.05 V 2+ Sn  /Ag 3 Sn  -0.05 ...
Electrolysis and Mass Changes charge (Coulombs) = current (Amperes) x time (sec) 1 mole e -  = 96,500 C = 1 Faraday 19.8
Anode:   Cathode: = 0.0126 mol Ca = 0.50 g Ca 19.8 How much Ca will be produced in an electrolytic cell of molten CaCl 2  ...
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Electrochemistry

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Electrochemistry

  1. 1. Electrochemistry Chapter 19
  2. 2. Electron Transfer Reactions <ul><li>Electron transfer reactions are oxidation-reduction or redox reactions. </li></ul><ul><li>Results in the generation of an electric current (electricity) or be caused by imposing an electric current. </li></ul><ul><li>Therefore, this field of chemistry is often called ELECTROCHEMISTRY. </li></ul>
  3. 3. Oxidation half-reaction (lose e - ) Reduction half-reaction (gain e - ) 19.1 <ul><li>Electrochemical processes are oxidation-reduction reactions in which: </li></ul><ul><li>the energy released by a spontaneous reaction is converted to electricity or </li></ul><ul><li>electrical energy is used to cause a nonspontaneous reaction to occur </li></ul>0 0 2+ 2- 2Mg ( s ) + O 2 ( g ) 2MgO ( s ) 2Mg 2Mg 2+ + 4e - O 2 + 4e - 2O 2-
  4. 4. Terminology for Redox Reactions <ul><li>OXIDATION —loss of electron(s) by a species; increase in oxidation number; increase in oxygen. </li></ul><ul><li>REDUCTION —gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen. </li></ul><ul><li>OXIDIZING AGENT —electron acceptor; species is reduced. (an agent facilitates something; ex. Travel agents don’t travel, they facilitate travel) </li></ul><ul><li>REDUCING AGENT —electron donor; species is oxidized. </li></ul>
  5. 5. You can’t have one… without the other! <ul><li>Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons. </li></ul><ul><li>You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation </li></ul>LEO the lion says GER ! ose lectrons xidation ain lectrons eduction GER!
  6. 6. Another way to remember <ul><li>OIL RIG </li></ul>xidation s ose eduction s ain
  7. 7. Review of Oxidation numbers The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. <ul><li>Free elements (uncombined state) have an oxidation number of zero. </li></ul>Na, Be, K, Pb, H 2 , O 2 , P 4 = 0 <ul><li>In monatomic ions, the oxidation number is equal to the charge on the ion. </li></ul>Li + , Li = +1 ; Fe 3+ , Fe = +3 ; O 2- , O = -2 <ul><li>The oxidation number of oxygen is usually –2 . In H 2 O 2 and O 2 2- it is –1 . </li></ul>4.4
  8. 8. <ul><li>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 . </li></ul>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. <ul><li>Group IA metals are +1 , IIA metals are +2 and fluorine is always –1 . </li></ul>HCO 3 - O = -2 H = +1 3x( -2 ) + 1 + ? = -1 C = +4 4.4 Oxidation numbers of all the atoms in HCO 3 - ?
  9. 9. Balancing Redox Equations 19.1 <ul><li>Write the unbalanced equation for the reaction ion ionic form. </li></ul>The oxidation of Fe 2+ to Fe 3+ by Cr 2 O 7 2- in acid solution? <ul><li>Separate the equation into two half-reactions. </li></ul>Oxidation: Reduction: <ul><li>Balance the atoms other than O and H in each half-reaction. </li></ul>Fe 2+ + Cr 2 O 7 2- Fe 3+ + Cr 3+ Cr 2 O 7 2- Cr 3+ +6 +3 Fe 2+ Fe 3+ +2 +3 Cr 2 O 7 2- 2Cr 3+
  10. 10. Balancing Redox Equations <ul><li>For reactions in acid, add H 2 O to balance O atoms and H + to balance H atoms. </li></ul><ul><li>Add electrons to one side of each half-reaction to balance the charges on the half-reaction. </li></ul><ul><li>If necessary, equalize the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients. </li></ul>19.1 Cr 2 O 7 2- 2Cr 3+ + 7H 2 O 14H + + Cr 2 O 7 2- 2Cr 3+ + 7H 2 O Fe 2+ Fe 3+ + 1e - 6e - + 14H + + Cr 2 O 7 2- 2Cr 3+ + 7H 2 O 6Fe 2+ 6Fe 3+ + 6e - 6e - + 14H + + Cr 2 O 7 2- 2Cr 3+ + 7H 2 O
  11. 11. Balancing Redox Equations <ul><li>Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel. You should also cancel like species. </li></ul>Oxidation: Reduction: <ul><li>Verify that the number of atoms and the charges are balanced . </li></ul>14x1 – 2 + 6x2 = 24 = 6x3 + 2x3 19.1 <ul><li>For reactions in basic solutions, add OH - to both sides of the equation for every H + that appears in the final equation. You should combine H + and OH - to make H 2 O. </li></ul>6e - + 14H + + Cr 2 O 7 2- 2Cr 3+ + 7H 2 O 6Fe 2+ 6Fe 3+ + 6e - 14H + + Cr 2 O 7 2- + 6Fe 2+ 6Fe 3+ + 2Cr 3+ + 7H 2 O
  12. 12. <ul><li>To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. </li></ul>CHEMICAL CHANGE ---> ELECTRIC CURRENT This is accomplished in a GALVANIC or VOLTAIC cell. A group of such cells is called a battery . http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf
  13. 13. Galvanic Cells 19.2 spontaneous redox reaction anode oxidation cat hode red uction - +
  14. 14. Galvanic Cells 19.2 <ul><li>The difference in electrical potential between the anode and cathode is called: </li></ul><ul><li>cell voltage </li></ul><ul><li>electromotive force (emf) </li></ul><ul><li>cell potential </li></ul>Cell Diagram [Cu 2+ ] = 1 M & [Zn 2+ ] = 1 M Zn ( s ) | Zn 2+ (1 M ) || Cu 2+ (1 M ) | Cu ( s ) anode cathode Zn ( s ) + Cu 2+ ( aq ) Cu ( s ) + Zn 2+ ( aq )
  15. 15. Standard Electrode Potentials 19.3 Zn ( s ) | Zn 2+ (1 M ) || H + (1 M ) | H 2 (1 atm) | Pt ( s ) Anode (oxidation): Cathode (reduction): 2e - + 2H + (1 M ) H 2 (1 atm ) Zn ( s ) Zn 2+ (1 M ) + 2e - Zn ( s ) + 2H + (1 M ) Zn 2+ + H 2 (1 atm )
  16. 16. Standard Electrode Potentials 19.3 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. E 0 = 0 V Standard hydrogen electrode (SHE) Reduction Reaction 2e - + 2H + (1 M ) H 2 (1 atm )
  17. 17. 19.3 <ul><li>E 0 is for the reaction as written </li></ul><ul><li>The more positive E 0 the greater the tendency for the substance to be reduced </li></ul><ul><li>The half-cell reactions are reversible </li></ul><ul><li>The sign of E 0 changes when the reaction is reversed </li></ul><ul><li>Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E 0 </li></ul>
  18. 18. 19.3 2 Standard Electrode Potentials Zn ( s ) | Zn 2+ (1 M ) || H + (1 M ) | H 2 (1 atm) | Pt ( s ) If the reaction is backwards, be sure to flip the sign! +2 So E o Zn/Zn = + 0.76 V +2 E 0 = 0 + 0.76 V = 0.76 V cell Standard emf (E 0 ) cell Zn 2+ (1 M ) + 2e - Zn E 0 = -0.76 V E 0 = E H /H + E Zn /Zn cell 0 0 + E 0 = E cathode + E anode cell 0 0
  19. 19. Standard Electrode Potentials 19.3 Pt ( s ) | H 2 (1 atm ) | H + (1 M ) || Cu 2+ (1 M ) | Cu ( s ) Anode (oxidation): Cathode (reduction): E cell = E Cu /Cu + E H /H+ 2+ 2 0 0 0 2e - + Cu 2+ (1 M ) Cu (s) H 2 (1 atm ) 2H + (1 M ) + 2e - H 2 (1 atm ) + Cu 2+ (1 M ) Cu ( s ) + 2H + (1 M ) E 0 = E cathode + E anode cell 0 0 E 0 = 0.34 V cell 0.34 = E Cu /Cu + - 0 0 2+ E Cu /Cu = 0.34 V 2+ 0
  20. 20. What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 M Cd(NO 3 ) 2 solution and a Cr electrode in a 1.0 M Cr(NO 3 ) 3 solution? Cd is the stronger oxidizer Cd will oxidize Cr Anode (oxidation): Cathode (reduction): x 2 x 3 19.3 Cd 2+ ( aq ) + 2e - Cd ( s ) E 0 = -0.40 V Cr 3+ ( aq ) + 3e - Cr ( s ) E 0 = -0.74 V 2e - + Cd 2+ (1 M ) Cd (s) Cr (s) Cr 3+ (1 M ) + 3e - 2Cr ( s ) + 3Cd 2+ (1 M ) 3Cd ( s ) + 2Cr 3+ (1 M ) E 0 = E cathode + E anode cell 0 0 E 0 = -0.40 + (+0.74) cell E 0 = 0.34 V cell
  21. 21. 19.4 Spontaneity of Redox Reactions  G = - nFE cell n = number of moles of electrons in reaction = 96,500 C/mol  G 0 = - RT ln K  G 0 = - nFE cell 0 F = 96,500 J V • mol = - nFE cell 0 E cell 0 = RT nF ln K (8.314 J/K • mol)(298 K) n (96,500 J/V • mol) ln K = = 0.0257 V n ln K E cell 0 = 0.0592 V n log K E cell 0
  22. 22. Spontaneity of Redox Reactions 19.4
  23. 23. Oxidation: Reduction: 19.4 n = 2 K = 1.23 x 10 -42 0 0 2+ + 2e - + Fe 2+ Fe 2Ag 2Ag + + 2e - What is the equilibrium constant for the following reaction at 25 0 C? Fe 2+ ( aq ) + 2Ag ( s ) Fe ( s ) + 2Ag + ( aq ) = 0.0257 V n ln K E cell 0 E 0 = -0.44 + -0.80 E 0 = -1.24 V 0.0257 V x n E 0 cell exp K = 0.0257 V x 2 -1.24 V = exp E 0 = E Fe /Fe + E Ag /Ag
  24. 24. The Effect of Concentration on Cell Emf  G =  G 0 + RT ln Q  G = - nFE - nFE = - nFE 0 + RT ln Q Nernst equation At 298 19.5  G 0 = - nFE 0 E = E 0 - ln Q RT nF - 0.0257 V n ln Q E 0 E = - 0.0592 V n log Q E 0 E =
  25. 25. Oxidation: Reduction: n = 2 E 0 = E Fe /Fe + E Cd /Cd 0 0 2+ 2+ E = 0.013 E > 0 Spontaneous 19.5 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 - E 0 = -0.44 + -(-0.40) E 0 = -0.04 V - 0.0257 V n ln Q E 0 E = - 0.0257 V 2 ln -0.04 V E = 0.010 0.60
  26. 26. Charging a Battery When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal. In your car, the battery charger is called an alternator. If you have a dead battery, it could be the battery needs to be replaced OR the alternator is not charging the battery properly.
  27. 27. Batteries 19.6 Leclanch é cell Dry cell Anode: Cathode: Zn (s) Zn 2+ ( aq ) + 2e - 2NH 4 ( aq ) + 2MnO 2 ( s ) + 2e - Mn 2 O 3 ( s ) + 2NH 3 ( aq ) + H 2 O ( l ) + Zn ( s ) + 2NH 4 ( aq ) + 2MnO 2 ( s ) Zn 2+ ( aq ) + 2NH 3 ( aq ) + H 2 O ( l ) + Mn 2 O 3 ( s )
  28. 28. Batteries Anode: Cathode: Mercury Battery 19.6 Zn(Hg) + 2OH - ( aq ) ZnO ( s ) + H 2 O ( l ) + 2e - HgO ( s ) + H 2 O ( l ) + 2e - Hg ( l ) + 2OH - ( aq ) Zn(Hg) + HgO ( s ) ZnO ( s ) + Hg ( l )
  29. 29. Batteries 19.6 Anode: Cathode: Lead storage battery PbO 2 ( s ) + 4H + ( aq ) + SO 2- (aq ) + 2e - PbSO 4 ( s ) + 2H 2 O ( l ) 4 Pb (s) + SO 2- ( aq ) PbSO 4 ( s ) + 2e - 4 Pb ( s ) + PbO 2 ( s ) + 4H + ( aq ) + 2SO 2- ( aq ) 2PbSO 4 ( s ) + 2H 2 O ( l ) 4
  30. 30. Batteries 19.6 Solid State Lithium Battery
  31. 31. Batteries 19.6 A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning Anode: Cathode: O 2 ( g ) + 2H 2 O ( l ) + 4e - 4OH - ( aq ) 2H 2 ( g ) + 4OH - ( aq ) 4H 2 O ( l ) + 4e - 2H 2 ( g ) + O 2 ( g ) 2H 2 O ( l )
  32. 32. Corrosion 19.7
  33. 33. Cathodic Protection of an Iron Storage Tank 19.7
  34. 34. 19.8 Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur.
  35. 35. Electrolysis of Water 19.8
  36. 36. Chemistry In Action: Dental Filling Discomfort Hg 2 /Ag 2 Hg 3 0.85 V 2+ Sn /Ag 3 Sn -0.05 V 2+ Sn /Ag 3 Sn -0.05 V 2+
  37. 37. Electrolysis and Mass Changes charge (Coulombs) = current (Amperes) x time (sec) 1 mole e - = 96,500 C = 1 Faraday 19.8
  38. 38. Anode: Cathode: = 0.0126 mol Ca = 0.50 g Ca 19.8 How much Ca will be produced in an electrolytic cell of molten CaCl 2 if a current of 0.452 A is passed through the cell for 1.5 hours? Ca 2+ ( l ) + 2e - Ca ( s ) 2Cl - ( l ) Cl 2 ( g ) + 2e - Ca 2+ ( l ) + 2Cl - ( l ) Ca ( s ) + Cl 2 ( g ) 2 mole e - = 1 mole Ca mol Ca = 0.452 C s x 1.5 hr x 3600 s hr 96,500 C 1 mol e - x 2 mol e - 1 mol Ca x

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