CH2205 - Electroanalytical Techniques

Electrochemical Cells Typically a cell is galvanic if it produces electrical energy or electrolytic if it consumes electrical energy. A cell is made of two conductors, or , each immersed in a suitable electrolyte. For electricity to flow it is necessary that the electrodes are connected externally by a metal conductor and that the two electrolyte solutions are in contact to permit movement of from one solution to the other.

Two common cell set-ups are;

V V Salt Bridge

Frit

Electricity is conducted in various ways in the cell. There is a migration of cations and anions as electrons move from one to the other, at the electrode surfaces an oxidation or reduction process provides a mechanism by which the ionic conduction of the solution is coupled with the electron conduction of the electrodes, providing a complete circuit for the flow of electricity.

By definition; An anode is where oxidation occurs. A cathode is where reduction occurs. The liquid junction (frit or salt bridge) is used to avoid the direct reaction of the components of the two half-cells, there would be a direct deposition of one metal onto the other. There is a small potential, the junction potential, arises at the interface between two electrolytic solutions of differing compositions.

Schematic representations of cells are often used to simplify the diagrams using the following notation;

2+ 2+ M(s) І M (aq)||Me (aq) І Me(s)

Where each line represents a phase boundary at which a potential may develop, two lines (||) represents a salt bridge and a long single line (|) represents a frit.

Cell Potentials Using the example cell;

- + 2AgCl(s)+H2(g)⇋2Ag(s)+2Cl +2H

The equilibrium constant for this reaction is; [H+]2[Cl−]2 K = p(H2)

To consider the value at any point during the reaction (using instantaneous concentration values) using Q;

+ 2 − 2 [H ]a [Cl ]a 푄 = p(H2)a

The change of the free energy in the cell is given by;

∆G = ∆Geq − ∆Ga

∆G = −RTlnK and ∆Ga = −RTlnQ ∴ ∆G = RTlnQ − RTlnK

The magnitude of free energy for the system depends on how far away the system is from equilibrium.

Since;

∆G = −nFEcell −nFEcell = RTlnQ − RTlnK −RTlnQ RT E = + lnK cell nF nF RT lnK = Eo nF cell RT ∴ E = E − lnQ cell o nF

Electrode Potentials A cell is made of two half-cell reactions, conventionally written as reductions (a species gaining electrons).

To obtain the cell, the second is subtracted from the first to cancel out electrons. To obtain the cell potential, strictly speaking, it must be done via free energy of the system, however if n is the same for each half-cell;

Ecell=Ecathode-Eanode

Calculating Half-Cell Potentials It is rare that a cell will be ‘standard’ and its potential will often have to be calculated via the . For the reaction; pP+qQ+ne-⇋rR+sS Then; RT [R]r[S]s E = E0 − ln nF [P]p[Q]q

The standard potential is often defined as the electrode potential of a half-cell when all reactants and products exist at unit activity. It is a physical constant that gives a quantitative description of the relative driving force of a half-cell reaction.

Potentiometry Potentiometry has been used for a long time in detecting the end-point during titrimetric analysis. More recently it has been used for the quantitative analysis of specific ions in solution. The method simply involves the determination of the potential between two half-cells, if one half- cell is constant then the potential of the other is effectively being measured and this potential is related to the concentration through the Nernst equation.

The equipment required is similar to that of the above cel