Electrochemistry: The branch of chemistry that covers the relative strengths of oxidants and reductants, the production of electric current from chemical reactions, and the use of electricity to produce chemical change.

Electrochemical cell: A system made up of two electrodes in contact with an electrolyte.

Electrode: A conductor of electricity, commonly a metal or graphite in contact with an electrolyte in an electrochemical cell.

Electrode reaction: A chemical reaction occurring at an electrode involving gain or loss of electrons. It is called a half-reaction. (See section 12d)
Examples:
Cu²⁺+ 2e¯ Cu; Zn Zn²⁺ + 2e¯

Redox couple: The two species of a half- reaction involving oxidation or reduction.
(See section 12d) Represented as oxidised species/reduced species
Examples:
Cu²⁺/Cu; Cl₂/Cl¯; Fe³⁺/Fe²⁺.

Half-cell: An electrode and the couple it is in contact with
Examples:

Zn(s) |Zn²⁺(aq)
a zinc rod in contact with an aqueous solution containing zinc cations
Fe³⁺(aq),Fe²⁺(aq)|Pt
a platinum rod in contact with an aqueous solution containing ferric and ferrous ions

The electrode may be one of the species of the couple or an inert species as with zinc and platinum respectively above. The vertical line represents a phase boundary.

Galvanic cell: an electrochemical cell that produces electricity from a chemical reaction.
It consists of two half-cells connected by a junction [e.g. a membrane (diaphragm) or salt bridge] allowing ions to be transferred between the electrolytes of the two half-cells. When the electrodes are connected externally by a conductor electrons flow through the conductor from the negative electrode to the positive electrode.
It is normally represented by two half-cells separated by a double vertical line representing the junction allowing transfer of ions.

how to represent a cell

Cell potential or electromotive force: symbol E, the electric potential difference between the electrodes of a galvanic cell when no current is flowing. In the above representation of a galvanic cell it is the electrode potential of the right hand electrode minus the electrode potential of the left hand electrode. It is normal to show the negative electrode on the left and the positive electrode on the right and thus express the cell potential as a positive quantity.
The SI unit of E is the volt, symbol V.

Standard hydrogen electrode: symbol she, a reference electrode defined as having zero electrode potential. It consists of a platinum electrode in contact with gaseous dihydrogen and aqueous hydronium ions under defined conditions.

Electrode potential (of a couple): symbol E(oxidised form/reduced form), also called reduction potential or redox potential is the electrode potential of that couple relative to the standard hydrogen electrode. More precisely it is the cell potential of a galvanic cell in which one half-cell is the standard hydrogen electrode, and is negative if the electrode of the couple is the negative electrode of that cell and positive if the electrode is the positive electrode of that cell.
Examples:

E(Cu2+/Cu)	=	0.34 V
E(Zn2+/Zn)	=	−0.76 V
E(Cl2/Cl¯)	=	1.40 V

The magnitude of a redox potential is a measure of the relative strength of an oxidant or reductant. The more positive the redox potential the stronger the oxidant (i.e. the greater its potential to oxidise some other species) and the weaker the reductant. Conversely the more negative the redox potential the stronger the reductant and the weaker the oxidant.

Table of electrode potentials: a table which lists the electrode potential of redox couples

Couple	Half-reaction	E / V
Zn2+/Zn	Zn2+ + 2e¯ Zn	−0.76
Sn2+/Sn	Sn2+ + 2e¯ Sn	−0.14
Cu2+/Cu	Cu2+ + 2e¯ Cu	0.34
I2/I¯	I2 + 2e¯ 2I¯	0.54
Fe3+/Fe2+	Fe3+ + e¯ Fe2+	0.77
Ag+/Ag	Ag+ + e¯ Ag	0.80
Br2/Br¯	Br2 + 2e¯ 2Br¯	1.10
O2/H2O	O2 + 4H+ + 4e¯ 2H2O	1.23
Cl2/Cl¯	Cl2 + 2e¯ 2Cl¯	1.40

Strong oxidant of a redox couple: An oxidant with a large positive redox potential
Examples
Fluorine, F₂, E(F₂/F¯) = 2.9 V;
The half-reaction is:
F₂ + 2e¯ 2F¯
Dichromate, Cr₂O₇²¯, E(Cr₂O₇/Cr³⁺) = 1.4 V.
The half-reaction is:
14H⁺ + Cr₂O₇²¯ + 6e¯ 2Cr³⁺ + 7H₂O .

Strong reductant of a redox couple: A reductant with a large negative redox potential
Example:
Sodium, Na, E(Na⁺/Na) = -2.7 V.
The half-reaction is:
Na Na⁺ + e¯
Magnesium, Mg, E(Mg²⁺/Mg) = -2.36 V
The half reaction is
Mg Mg²⁺ + 2e¯

As electrons flow spontaneously from the negative to the positive electrode the electode potentials can be used to deduce the direction of a redox reaction.
Example:
Can Cl₂ oxidise copper, Cu?
Since this reaction would involve flow of electrons from Cu to Cl₂ it can be seen from the electrode potentials above that the redox reaction can occur as E(Cu²⁺/Cu) is negative with respect to E(Cl₂/Cl¯).
A table of electrode potentials in which the electrode potentials increase down the table (most negative at the top and most positive at the bottom) can be regarded as an energy diagram. A spontaneous reaction involving two couples will be that in which electrons lose energy by moving from a reductant of one couple to the oxidant of another couple below it in the table.

Examples:

Couple	E (V)
Zn2+ / Zn	−0.76
Cu2+ / Cu	0.34
Cl2 / Cl¯	1.40

Electrons can flow from Zn to Cu²⁺ but not from Cu to Zn²⁺ Thus zinc metal can reduce copper ions to copper metal but copper metal cannot reduce zinc ions to zinc metal.

Predict whether:

4. Sn can reduce Br₂
Answer 4

5. O₂ can oxidise Cl¯
Answer 5

6. Ag can reduce Cu²⁺
Answer 6

7. Fe³⁺ can oxidise I¯
Answer 7

For the cell
Zn²⁺|Zn||Cu²⁺|Cu
the half reactions will be:
Zn Zn²⁺ + 2e¯ (as the zinc elelectrode is the more negative)
and
Cu²⁺ + 2e¯ Cu (as the copper electrode is the more positive)
and the overall reaction the sum of these:
Zn + Cu²⁺ Zn²⁺ + Cu

The cell potential is the difference of the two electrode potentials, i.e. the potential of the positive electrode minus the potential of the negative electrode.
E(cell) = E(Cu²⁺/Cu) - E(Zn²⁺/Zn) = 0.34 V - (-0.76 V) = 1.10 V

WARNING

Write the half reaction equations, the chemical equation for the overall cell reaction
and determine the cell potential for a galvanic cell made from the redox couples:

8. Zn²⁺/Zn and Fe³⁺/Fe²⁺
Answer 8

the "lemon cell"

the vanadium cell reaction