Electrochemistry

1
Electrochemistry

• Applications of Redox

2
Review

• Oxidation reduction reactions involve a transfer
  of electrons.
• OIL- RIG
• Oxidation Involves Gain
• Reduction Involves Loss
• LEO-GER
• Lose Electrons Oxidation
• Gain Electrons Reduction

3
Applications

• Moving electrons is electric current.
• 8HMnO4- 5Fe2 5e- Mn2 5Fe3
  4H2O
• Helps to break the reactions into half reactions.
• 8HMnO4-5e- Mn2 4H2O
• 5(Fe2 Fe3 e- )
• In the same mixture it happens without doing
  useful work, but if separate

4

• Connected this way the reaction starts
• Stops immediately because charge builds up.

H MnO4-
Fe2
5
Galvanic Cell
Salt Bridge allows current to flow
H MnO4-
Fe2
6

• Electricity travels in a complete circuit
• Instead of a salt bridge

H MnO4-
Fe2
7
Porous Disk
H MnO4-
Fe2
8
e-
e-
e-
e-
Anode
Cathode
e-
e-
Reducing Agent
Oxidizing Agent
9
Cell Potential

• Oxidizing agent pushes the electron.
• Reducing agent pulls the electron.
• The push or pull (driving force) is called the
  cell potential Ecell
• Also called the electromotive force (emf)
• Unit is the volt(V)
• 1 joule of work/coulomb of charge
• Measured with a voltmeter

10
0.76
H2 in
Cathode
Anode
H Cl-
Zn2 SO4-2
1 M HCl
1 M ZnSO4
11
Standard Hydrogen Electrode

• This is the reference all other oxidations are
  compared to
• Eº 0
• º indicates standard states of 25ºC, 1 atm, 1
  M solutions.

H2 in
H Cl-
1 M HCl
12
Cell Potential

• Zn(s) Cu2 (aq) Zn2(aq) Cu(s)
• The total cell potential is the sum of the
  potential at each electrode.
• Eº cell EºZn Zn2 Eº Cu2 Cu
• We can look up reduction potentials in a table.
• One of the reactions must be reversed, so change
  it sign.

13
Cell Potential

• Determine the cell potential for a galvanic cell
  based on the redox reaction.
• Cu(s) Fe3(aq) Cu2(aq) Fe2(aq)
• Fe3(aq) e- Fe2(aq) Eº 0.77 V
• Cu2(aq)2e- Cu(s) Eº 0.34 V
• Cu(s) Cu2(aq)2e- Eº -0.34 V
• 2Fe3(aq) 2e- 2Fe2(aq) Eº 0.77 V

14
Line Notation

• solid½Aqueous½½Aqueous½solid
• Anode on the left½½Cathode on the right
• Single line different phases.
• Double line porous disk or salt bridge.
• If all the substances on one side are aqueous, a
  platinum electrode is indicated.
• For the last reaction
• Cu(s)½Cu2(aq)½½Fe2(aq),Fe3(aq)½Pt(s)

15
Galvanic Cell

• The reaction always runs spontaneously in the
  direction that produced a positive cell
  potential.
• Four things for a complete description.
• Cell Potential
• Direction of flow
• Designation of anode and cathode
• Nature of all the components- electrodes and ions

16
Practice

• Completely describe the galvanic cell based on
  the following half-reactions under standard
  conditions.
• MnO4- 8 H 5e- Mn2 4H2O Eº1.51
• Fe3 3e- Fe(s) Eº0.036V

17
Potential, Work and DG

• emf potential (V) work (J) / Charge(C)
• E work done by system / charge
• E -w/q
• Charge is measured in coulombs.
• -w qE
• Faraday 96,485 C/mol e-
• q nF moles of e- x charge/mole e-
• w -qE -nFE DG

18
Potential, Work and DG

• DGº -nFE º
• if E º lt 0, then DGº gt 0 spontaneous
• if E º gt 0, then DGº lt 0 nonspontaneous
• In fact, reverse is spontaneous.
• Calculate DGº for the following reaction
• Cu2(aq) Fe(s) Cu(s) Fe2(aq)
• Fe2(aq) e- Fe(s) Eº 0.44 V
• Cu2(aq)2e- Cu(s) Eº 0.34 V

19
Cell Potential and Concentration

• Qualitatively - Can predict direction of change
  in E from LeChâtelier.
• 2Al(s) 3Mn2(aq) 2Al3(aq) 3Mn(s)
• Predict if Ecell will be greater or less than
  Eºcell if Al3 1.5 M and Mn2 1.0 M
• if Al3 1.0 M and Mn2 1.5M
• if Al3 1.5 M and Mn2 1.5 M

20
The Nernst Equation

• DG DGº RTln(Q)
• -nFE -nFEº RTln(Q)
• E Eº - RTln(Q) nF
• 2Al(s) 3Mn2(aq) 2Al3(aq) 3Mn(s) Eº
  0.48 V
• Always have to figure out n by balancing.
• If concentration can gives voltage, then from
  voltage we can tell concentration.

21
The Nernst Equation

• As reactions proceed concentrations of products
  increase and reactants decrease.
• Reach equilibrium where Q K and Ecell 0
• 0 Eº - RTln(K) nF
• Eº RTln(K) nF
• nFEº ln(K) RT

22
Batteries are Galvanic Cells

• Car batteries are lead storage batteries.
• Pb PbO2 H2SO4 PbSO4(s) H2O
• Dry Cell Zn NH4 MnO2 Zn2 NH3
  H2O
• Alkaline Zn MnO2 ZnO Mn2O3 (in base)
• NiCad
• NiO2 Cd 2H2O Cd(OH)2 Ni(OH)2

23
Corrosion

• Rusting - spontaneous oxidation.
• Most structural metals have reduction potentials
  that are less positive than O2 .
• Fe Fe2 2e- Eº 0.44 V
• O2 2H2O 4e- 4OH- Eº 0.40 V
• Fe2 O2 H2O Fe2 O3 H
• Reaction happens in two places.

24
Salt speeds up process by increasing conductivity
Water
Rust
e-
Iron Dissolves- Fe Fe2
25
Preventing Corrosion

• Coating to keep out air and water.
• Galvanizing - Putting on a zinc coat
• Has a lower reduction potential, so it is more.
  easily oxidized.
• Alloying with metals that form oxide coats.
• Cathodic Protection - Attaching large pieces of
  an active metal like magnesium that get oxidized
  instead.

26
Electrolysis

• Running a galvanic cell backwards.
• Put a voltage bigger than the potential and
  reverse the direction of the redox reaction.
• Used for electroplating.

27
1.10
e-
e-
Zn
Cu
1.0 M Cu2
1.0 M Zn2
Cathode
Anode
28
A battery gt1.10V
e-
e-
Zn
Cu
1.0 M Cu2
1.0 M Zn2
Cathode
Anode
29
Calculating plating

• Have to count charge.
• Measure current I (in amperes)
• 1 amp 1 coulomb of charge per second
• q I x t
• q/nF moles of metal
• Mass of plated metal
• How long must 5.00 amp current be applied to
  produce 15.5 g of Ag from Ag

30
Other uses

• Electroysis of water.
• Seperating mixtures of ions.
• More positive reduction potential means the
  reaction proceeds forward.
• We want the reverse.
• Most negative reduction potential is easiest to
  plate out of solution.