Introduction to Electrochemistry

1

• Introduction to Electrochemistry
• Terms
• Oxidation is the loss of electrons by a substance
• Cu Cu2 2e-
• 2H2O 4H O2(g) 4e-
• Reduction is the gain of electrons by a substance
• Zn2 2e- Zn(s)
• 2H 2e- H2(g)
• Oxidation cannot occur in the absence of
  reduction
• In a reaction in which oxidation-reduction occurs
  the number of electrons gained by the substance
  reduced must equal the number of electrons lost
  by the substance oxidized
• This principle is used in balancing
  oxidation-reduction equations
• Convenient formalism - separate the overall
  equation into two half- equations, one an
  oxidation half-equation, the other a reduction
  half-equation
• Cu2 Zn Cu Zn2
• Cu2 2e- Cu
• Zn Zn2 2e-

2

• Introduction to Electrochemistry
• Oxidation of Fe2 with Cr2O72- in acid
• 6x(Fe2
  Fe3 e-)
• Cr2O72- 14H 6e- 2Cr3 7H2O
• Cr2O72- 14H 6Fe2 2Cr3 6Fe3
  7H2O
• These kinds of reactions can be carried out in
  homogeneous solution, in heterogeneous mixtures
  or at the surface of inert electrical
  conductors
• Electrochemical cells - example Daniel Cell
  involving the Zn/Cu2 couple
• The Daniel cell is an example

• Current flows when two electrodes connected with
  a conductor
• Zn is consumed and Cu is precipitated
• Equilibrium will be established
• Zn Cu2 Zn2 Cu
• Ratio Zn2/Cu2 becomes constant and very
  large
• The same ratio would be established if Zn were
  inserted into a Cu2 solution

3

• Introduction to Electrochemistry
• Electrochemical cells
• If the reactants are not at equilibrium the
  tendency to reach equilibrium drives electrons
  through the connecting wires and ions through the
  solution
• This tendency is represented by a voltage or EMF
  which is related to the equilibrium constant
• Galvanic or voltaic cells produce energy as the
  reaction approaches equilibrium
• Electrolytic cells consume electricity provided
  by an external source to drive the reaction away
  from the approach towards equilibrium
• Reversible cells produce exactly the reverse
  reaction in the electrolytic direction as the
  galvanic direction
• Current passage in cells
• In the external cells electrons carry the charge
• In the internal part of the cell ions carry the
  charge
• For the Daniel cell
• Zn electrode is negative since it is the source
  of electrons and Zn2 is produced in the
  solution since Zn is oxidized
• This means negative ions must move into the Zn
  compartment to maintain charge balance
• Postive ions migrate out of the Zn compartment
• Cu2 is consumed and positive ions move to the Cu
  compartment to maintain charge balance

4

• Introduction to Electrochemistry
• Naming electrodes
• Anode the electrode at which oxidation occurs
• Cathode the electrode at which reduction occurs
• Electrode potentials

• Electrode potentials are determined relative to
  an electrode called the standard hydrogen
  electrode-SHE
• 2H 2e- H2(g) Eo 0.0000 V at all temp
• If the potential of the Zn electrode is measured
  with respect to SHE, E -0.76 V
• Zn is negative and thus is the source of
  electrons and must be the anode in a galvanic
  cell
• Zn Zn2 2e- anode
• 2H 2e- H2(g) cathode
• 2H Zn Zn2 H2(g)
• If Cu replaces Zn and Cu2 replaces Zn2, SHE
  becomes anode and Cu becomes cathode
• Cu2 is positive, SHE is negative

5

• Introduction to Electrochemistry
• IUPAC sign convention for standard electrode
  potentials
• Write the half-equation for the reaction of
  interest as a reduction
• Give the potential the sign of the polarity the
  electrode has when it is made part of a galvanic
  cell when the other electrode is SHE
• Examples
• Zn2 2e- Zn Eo -0.763 V
• Cu2 2e- Cu Eo 0.337 V
• Effect of concentration on electrode potential
• For the general equation
• aA bB ne- cC dD
• Eo standard electrode potential
• R universal gas constant 8.314 volt-coul K-1
  mol-1
• T absolute temperature
• nnumber of e- shown in the balanced equation
• ? Faradays constant 96,495.3 coul/mol
• ln natural log 2.303 log10

If T298K,
6

• Introcduction to Electrochemistry
• Effect of concentration on electrode potential
• Examples equations are written as reduction
  half-equations
• 2H 2e- H2(g)
• ?r2O72- 14H 6e- 2Cr3 7H2O
• AgCl 1e- Ag Cl-
• When Q 0, E Eo
• E is generally temperature dependent
• E is a relative quantity
• Gives the sign of electrode at which the
  specified reaction occurs when the electrode is
  a part of the cell in which SHE is the other
  electrode
• The E is independent of the stoichiometry written
  for the balanced equation
• 2H 2e- H2(g)
• H e- H2(g)
• Appendix 5 in FAC7 gives Eo for many
  half-equations sorted by element

7

• Introduction to Electrochemistry
• Example Calculate the potential of a piece of Pt
  inserted in a solution prepared by saturating a
  0.0100 M solution of KBr with Br2(l) vs. SHE
• Br2(l) 2e- 2Br-
• Schematic representation of electrochemical cells
• Write the anode components on the left
• Write the cathode components on the right
• Give the activity or concentrations of reagents
• Represent any phase interface with a vertical
  line
• Salt bridges are represented with two vertical
  lines
• Example the Daniel cell
• ZnZnSO4(aZn21)CuSO4(aCu21)Cu
• Cell potentials - Example Daniel cell
• Cathode Cu2 2e- Cu
  Eo 0.337 V
• Anode Zn Zn2 2e-
  Eo 0.763 V
• Overall Cu2 Zn Zn2 Cu
  Ecello 1.100 V ? Ecell gt 0, reaction
  spontaneous
• Galvanic cell

8
Cell potentials - Example CuCuSO4(aCu21)ZnSO
4(aZn21)Zn Cathode Zn2 2e- Zn
Eo -0.763 V Anode
Cu Cu2 2e- Eo -0.337
V Overall Cu2 Zn Zn2 Cu
Ecello -1.100 V ? Ecell lt 0, reaction not

spontaneous electrolytic
cell Example PtH2(g)(P1atm)HCl(aH0.5)Cu2
(aCu20.1)Cu Cathode Cu2 2e-
Cu Anode H2(g) 2H 2e-
9

• Introduction to Electrochemistry
• Equilibrium constants from cell potentials
• For oxidation-reduction equations from standard
  potentials
• aAox ne- aAred
• bBox ne- bBred
• aAox bBred aAred bBox
• if the system is at equilibrium EA EB
• for the Nernst equation

10

• Introduction to Electrochemistry
• Example calculate the equilibrium constant for
• Cr2O72- 14H 6e- 2Cr3 14H2O
• 6Fe2 6e-
  6Fe3
• Cr2O72- 6Fe2 14H 2Cr3 6Fe3
  14H2O

11

• Introduction to Electrochemistry
• Example A silver electrode immersed in a 1.00 x
  10-2 M solution of Na2SeO3 solution that is
  saturated with Ag2SeO3 acts as a cathode when
  coupled with an SHE. Calculate the Ksp for
  Ag2SeO4 if the cell potential is 0.450V
• Ksp Ag2SeO32-
• Cathode Ag e- Ag
• anode SHE