Chemistry and electricity:

The connection between chemistry and electricity is a very old one, going back to ALESSANDRO
VOLTA'S discovery, in 1793, that electricity could be produced by placing two dissimilar metals
on opposite sides of a moistened paper. In 1800, Nicholson and Carlisle, using Volta's primitive
battery as a source, showed that an electric current could decompose water into oxygen and
hydrogen. This was surely one of the most significant experiments in the history of chemistry, for
it implied that the atoms of hydrogen and oxygen were associated with positive and negative
electric charges, which must be the source of the bonding forces between them. By 1812, the
Swedish chemist BERZELIUS could propose that all atoms are electrified, hydrogen and the
metals being positive, the nonmetals negative. In electrolysis, the applied voltage was thought to
overpower the attraction between these opposite charges, pulling the electrified atoms apart in
the form of ions (named by Berzelius from the Greek for “travelers”). It would be almost exactly a
hundred years later before the shared electron pair theory of G.N. LEWIS could offer a significant
improvement over this view of chemical bonding.

Meanwhile the use of electricity as a means of bringing about chemical change continued to

play a central role in the development of chemistry. HUMPHREY DAVEY prepared the first
elemental sodium by electrolysis of a sodium hydroxide melt. It was left to Davey's former
assistant, MICHAEL FARADAY, to show that there is a direct relation between the amount of
electric charge passed through the solution and the quantity of electrolysis products. JAMES
CLERK MAXWELL immediately saw this as evidence for the “molecule of electricity”, but the
world would not be receptive to the concept of the electron until the end of the century.

Electrochemistry:

Electrochemistry is a branch of chemistry that studies chemical reactions which take place in

asolution at the interface of an electron conductor (a metal or a semiconductor) and an ionic conductor
(the electrolyte), and which involve electron transfer between the electrode and the electrolyte or species
in solution.

If a chemical reaction is driven by an external applied voltage, as in electrolysis, or if a voltage is created

by a chemical reaction as in a battery, it is an electrochemical reaction. In contrast, chemical reactions
where electrons are transferred between molecules are called oxidation/reduction (redox) reactions. In
general, electrochemistry deals with situations whereoxidation and reduction reactions are separated in
space or time, connected by an external electric circuit to understand each process.
Principles:

Redox Reactions:

Redox stands for reduction-oxidation, and are electrochemical processes involving electron transfer to

or from a molecule or ion changing its oxidation state. This reaction can occur through the application of
an external voltage or through the release of chemical energy.

Oxidation and reduction

xidation and reduction describe the change of oxidation state that takes place in the atoms, ions or
molecules involved in an electrochemical reaction. Formally, oxidation state is the
hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic.
An atom or ion that gives up an electron to another atom or ion has its oxidation state increase, and the
recipient of the negatively charged electron has its oxidation state decrease. Oxidation and reduction
always occur in a paired fashion such that one species is oxidized when another is reduced. This paired
electron transfer is called a redox reaction.

For example, when atomic sodium reacts with atomic chlorine, sodium donates one electron and attains
an oxidation state of +1. Chlorine accepts the electron and its oxidation state is reduced to −1. The sign of
the oxidation state (positive/negative) actually corresponds to the value of each ion's electronic charge.
The attraction of the differently charged sodium and chlorine ions is the reason they then form an ionic
bond.

The loss of electrons from an atom or molecule is called oxidation, and the gain of electrons is reduction.
This can be easily remembered through the use of mnemonic devices. Two of the most popular are "OIL
RIG" (Oxidation Is Loss, Reduction Is Gain) and "LEO" the lion says"GER" (Lose Electrons: Oxidization,
Gain Electrons: Reduction). For cases where electrons are shared (covalent bonds) between atoms with
large differences in electronegativity, the electron is assigned to the atom with the largest electronegativity
in determining the oxidation state.

The atom or molecule which loses electrons is known as the reducing agent, or reductant, and the
substance which accepts the electrons is called the oxidizing agent, or oxidant. The oxidizing agent is
always being reduced in a reaction; the reducing agent is always being oxidized. Oxygen is a common
oxidizing agent, but not the only one. Despite the name, an oxidation reaction does not necessarily need
to involve oxygen. In fact, a fire can be fed by an oxidant other than oxygen; fluorine fires are often
unquenchable, as fluorine is an even stronger oxidant (it has a higher electronegativity) than oxygen.

For reactions involving oxygen, the gain of oxygen implies the oxidation of the atom or molecule to which
the oxygen is added (and the oxygen is reduced). In organic compounds, such as butane or ethanol, the
loss of hydrogen implies oxidation of the molecule from which it is lost (and the hydrogen is reduced).
This follows because the hydrogen donates its electron in covalent bonds with non-metals but it takes the
electron along when it is lost. Conversely, loss of oxygen or gain of hydrogen implies reduction.

Balancing redox reactions

Electrochemical reactions in water are better understood by balancing redox reactions using the ion-
electron method where H+ , OH– ion,H2O and electrons (to compensate the oxidation changes) are added to
cell's half-reactions for oxidation and reduction.

Acidic medium

n acid medium H+ ions and water are added to half-reactions to balance the overall reaction. For example,
when manganese reacts withsodium bismuthate.

Unbalanced reaction: Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + MnO4–(aq)

Oxidation: 4 H2O(l) + Mn2+(aq) → MnO4–(aq) + 8 H+(aq) + 5 e–
Reduction: 2 e– + 6 H+(aq) + BiO3–(s) → Bi3+(aq) + 3 H2O(l)

Finally, the reaction is balanced by multiplying the number of electrons from the reduction half reaction to
oxidation half reaction and vice versa and adding both half reactions, thus solving the equation.

8 H2O(l) + 2 Mn2+(aq) → 2 MnO4–(aq) + 16 H+(aq) + 10 e–

10 e– + 30 H+(aq) + 5 BiO3–(s) → 5 Bi3+(aq) + 15 H2O(l)

Reaction balanced:

14 H+(aq) + 2 Mn2+(aq) + 5 NaBiO3(s) → 7 H2O(l) + 2 MnO4–(aq) + 5 Bi3+(aq) + 5 Na+(aq)

Basic medium

In basic medium OH– ions and water are added to half reactions to balance the overall reaction. For
example, on reaction between potassium permanganate and sodium sulfite.

Unbalanced reaction: KMnO4 + Na2SO3 + H2O → MnO2 + Na2SO4 + KOH

Reduction: 3 e– + 2 H2O + MnO4– → MnO2 + 4 OH–
Oxidation: 2 OH– + SO32– → SO42– + H2O + 2 e–

The same procedure as followed on acid medium by multiplying electrons to opposite half
reactions solve the equation thus balancing the overall reaction.

6 e– + 4 H2O + 2 MnO4– → 2 MnO2 + 8 OH–

6 OH– + 3 SO32– → 3 SO42– + 3 H2O + 6e–

Equation balanced:

2 KMnO4 + 3 Na2SO3 + H2O → 2 MnO2 + 3 Na2SO4 + 2 KOH

Neutral medium

he same procedure as used on acid medium is applied, for example on balancing using electron ion
method to complete combustion ofpropane.

Unbalanced reaction: C3H8 + O2 → CO2 + H2O

Reduction: 4 H+ + O2 + 4 e– → 2 H2O
Oxidation: 6 H2O + C3H8 → 3 CO2 + 20 e– + 20 H+

As in acid and basic medium, electrons which were used to compensate oxidation changes are multiplied
to opposite half reactions, thus solving the equation.

20 H+ + 5 O2 + 20 e– → 10 H2O

6 H2O + C3H8 → 3 CO2 + 20 e– + 20 H+

Equation balanced:

C3H8 + 5 O2 → 3 CO2 + 4 H2O

Electrochemical cells:
Although it is physically impossible to measure or manipulate the potential difference between a piece of
metal and the solution in which it is immersed, we can easily measure a potential difference between two
such electrodes immersed in a solution. The result will be the sum of the two electrode potentials, we
shall see farther on that such measurements can be supply all the information we need in order to
characterize the two electrode reactions.

Simple electrochemical cell

This arrangement is called a galvanic cell. A typical cell might consist of two pieces of metal, one zinc and
the other copper, each immersed each in a solution containing a dis-solved salt of the corresponding
metal. The two solutions are separated by a porous barrier that prevents them from rapidly mixing but
allows ions to diffuse through.
If we simply left it at that, no significant amount of reaction would take place. However, if we connect the
zinc and copper by means of a metallic conductor, the excess electrons that remain when Zn2+ ions go
into solution in the left cell would be able to flow through the external circuit and into the right electrode,
where they could be delivered to the Cu2+ ions which become “discharged”, that is, converted into Cu
atoms at the surface of the copper electrode. The net reaction is the same as before— the oxidation of
zinc by copper(II) ions:

but this time, the oxidation and reduction steps take place in separate locations:

Electrochemical cells allow measurement and control of a redox reaction:

The reaction can be started and stopped by connecting or disconnecting the two electrodes. If we place a
variable resistance in the circuit, we can even control the rate of the net cell reaction by simply turning a
knob. By connecting a battery or other source of current to the two electrodes, we can force the reaction
to proceed in its non-spontaneous, or reverse direction.