1. Electroanalytical Chemistry
Chapters 22, 23, 24 & 25
Electroanalytical Chemistry

2. Electroanalytical Chemistry

3. Electroanalytical Chemistry...
It encompasses a group of quantitative analytical methods that are based upon the electrical properties of a solution of the analyte when it is made part of an electrochemical cell.

4. Why Electroanalytical Chemistry?
Electroanalytical methods have certain advantages over other analytical methods. Electrochemical analysis allows for the determination of different oxidation states of an element in a solution, not just the total concentration of the element.
Electroanalytical techniques are capable of producing exceptionally low detection limits and an abundance of characterization information including chemical kinetics information. The other important advantage of this method is its low cost.

5. Galvanic Electrochemical Cell with Salt Bridge

6. Electroanalytical techniques are capable of producing exceptionally low detection limits and a wealth of characterization information describing electrochemically addressable systems. Such information includes stoichiometry and rate of interfacial charge transfer, rate of mass transfer, extent of adsorption or chemisorption, and rates and equilibrium constants for chemical reactions.

7. Oxidizing Agent + ne  Reduction
History
Polarography was first discovered by a czechoslovavian chemist by the name of Heyrovsky in He won the Nobel prize for it in He proposed that the current recording generated by a oxidation or reduction in a cell as the A.P. is continuously increased:
Oxidizing Agent + ne  Reduction

8. Oxygen Probe A.P. – 650 mV Ag|AgCl Reactions: O2 + 2H2O + 4e-  4OH-
4Ag + 4Cl-  4AgCl + 4e-

9. Hydrogen Probe A.P. +650 mV Reactions: H2O2  O2 +2H+ +2e-
2Ag+ + 2e-  2Ag

10. [Cu(OH2)6]2+(aq) + Zn --> Cu + [Zn(OH2)6]2+(aq)
Daniell Cell
This cell is based on the overall reaction
[Cu(OH2)6]2+(aq) + Zn --> Cu + [Zn(OH2)6]2+(aq)
and functions by dissolution of Zn from the anode and deposition of Cu at the cathode. It is therefore very simply represented as
Zn | [Zn(OH2)6]2+(aq) || [Cu(OH2)6]2+(aq) | Cu or just as Zn | Zn(II)(aq) || Cu(II)(aq) | Cu

11. Galvanic Cells
A galvanic cell consists of at least two half cells, a reduction cell and an oxidation cell. Chemical reactions in the two half cells provide the energy for the galvanic cell operations. The reactions always run spontaneously in the direction that produced a positive cell potential

12. Voltaic Cells
A voltaic cell is an electrochemical cell that external electrical current flow can be created using any two different metals since metals differ in their tendency to lose electrons. Zinc more readily electrons than copper, so placing zinc and copper metal in solutions of their salts can cause electrons to flow through an external wire which leads from the zinc to the copper. The following is a diagram of a voltaic cell.

13. Cathodes and Anodes

14. Reactions at the Anode Some examples are: Cu(s) <=> Cu2+ + 2e-
Fe2+ <=> Fe3+ + e-
H2(g) <=> 2H+ + 2e-
Ag(s) + Cl- <=> AgCl(s) + e-

15. Reactions at Cathodes
Electrons supplied by external circuit via an inert electrode (platinum or gold)
Some examples are:
Cu2+ + 2e- <=> Cu(s)
Fe3+ + e- <=> Fe2+
2H+ + 2e- <=> H2(g)
AgCl(s) + e- <=> Ag(s) + Cl-

16. Cells Without Liquid Junctions
Liquid junction - the interface between 2 different electrolytic solutions
Cell can contain more than one
A small Junction Potential arises at these interfaces
Sometimes it is possible to prepare cells that share a
Common electrolyte to avoid this problem

17. Concentration Cells
A concentration cell is an electrochemical cell in which the electrode couple at both electrodes is the same but the concentrations of substances at the two electrodes may differ. The potential difference across a concentration cell can be calculated using the Nernst equation.

18. Calculation of Cell Potential from Electrode Potentials
Nernst Equation: The cell potential for a voltaic cell under standard conditions can be calculated from the standard electrode potentials. But real voltaic cells will typically differ from the standard conditions. The Nernst equation relates the cell potential to its standard cell potential.

19. Cell Potential
One implication is that the cell potential will be reduced from the standard value if the concentration of Zn2+(aq) is greater than that of Cu2+(aq) at the standard temperature. An excess concentration of Cu2+(aq) will give a higher voltage. The graph at right shows the increase in cell voltage with increasing concentration of the cation. Note that the horizontal axis is logarithmic, and that the straight line variation of the voltage represents an logarithmic variation with Q. Note that the cell potential is equal to the standard value if the concentrations are equal even if they are not equal to the standard value of 1M, since the logarithm gives the value zero.

20. The Nernst Equation R = gas constant T = temperature in Kelvins
Q = thermodynamic reaction quotient
F = Faraday's constant
n = number of electrons transferred

21. Currents in Electrochemical Cells
Ohms law is usually obeyed:
E=IR
where E is the potential difference in volts responsible for the movement of the ions, I is the current in amps, and R is the resistance in ohms of the electrolyte to the current

27. Summary of Important Types of Polarography

29. LCEC...

36. The Ilkavic Equation:
The diffusion current, id, is in the plateau region of a polarogram and, thus, is independent of potential
id = 706n C D1/2 m2/3 t1/6
Where: id = max value of the diffusion current in the life of the drop
n = # of electrons involved
C = Cons. Substitution, mmoles/L
D = diffusion coefficient of the ion (cm2/s)
m = mass of Hg in mg/s
t = time between drops

37. References: http://www.anachem.umu.se/jumpstation.htm