Download Introduction to Electrochemistry

Chemistry 30 Notes
Introduction to Electrochemistry
During redox reactions, electrons pass from one substance to another. The flow of
electrons - current - can be harnessed to do work. Electrochemistry is the branch of
chemistry that deals with the conversion between chemical and electrical energy.
1. Electrochemical Cells
The basic unit of all batteries is the
electrochemical cell (also called a
galvanic cell). Electrochemical cells
convert the energy of a spontaneous
redox reaction into electricity.
An electrochemical cell consists of
two half-cells, each having an
electrode (solid electrical conductor)
in contact with a solution of ions.
The electrodes are where the redox
half-reactions occur.
1. Anode – is the site of oxidation
and the source of electrons,
making it the negative post of the
electrochemical cell
2. Cathode – site of reduction and
the positive post as it consumes
electrons
The electrodes are connected with wire. The force that moves the electrons from the
anode to the cathode via the wire is known as the electrical potential and it is measured
in volts, using a voltmeter. It is generally defined as the cell potential or voltage for a
source in a circuit.
Chemistry 30 Notes
Once the half-cells are connected, there is an immediate flow of electrons (current) but
then it abruptly stops as the charges build up in the half-cells. With nowhere to go, the
build-up of charge prevents further electron transfer and the redox reaction stops.
To solve this problem, the two half-cells are connected with a salt bridge, a U-shaped
tube filled with non-reactive electrolyte that is designed to keep ions flowing and the
cells electrically neutral. The ends of the tube are plugged with cotton balls to prevent
the solution from falling out but are porous enough to permit some fluid and ions to flow
between cells.
2. Electrochemical Cell Example
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
This reaction involves two half-reactions:
Zn → Zn +2 + 2 e-
Cu2+ + 2e- → Cu
oxidation
reduction
One half-cell contains a strip of zinc metal (which acts as an electrode)
in a Zn(NO3)2 solution. The other half-cell contains copper metal in
Cu(NO3)2. KNO3 is used as the salt bridge.
The zinc half-cell undergoes oxidation. Here, the solid zinc electrode
disintegrates, forming zinc ions and releasing electrons.
The copper half-cell undergoes reduction. Here, copper ions from the
electrolytic solution become deposited on the copper electrode,
forming more solid copper.
At the anode, Zn2+ ions are being produced and
go into solution. This causes a build-up of
positive ions in this solution. The excess
positive charge attracts the negative NO3ions from the salt bridge, thereby keeping the
solution electrically neutral.
Chemistry 30 Notes
At the cathode the opposite occurs. As positive Cu2+ ions are removed from solution, to
form solid Cu, the solution becomes overly negative. This attracts the positive K+ ions
from the salt bridge, keeping this side of the cell neutral.
External Circuit
Electrons flow
from A to C
Anode to Cathode
Internal Circuit
Anions
to the Anode
Cations
to the Cathode
As the reaction proceeds, the voltage will eventually decrease and finally become zero
when the concentrations of ions in solution reach equilibrium. The cell will then be
considered ‘dead’.
3. Predicting Cell Reactions and Equations
To identify the anode and cathode and
to calculate the voltage of the cell, we
will need to use a Table of Reduction
Potentials.
Standard Potentials (E0) are measured
against a hydrogen ion reduction
reaction, which is arbitrarily assigned a
potential of zero volts.
The cell’s total voltage (E0cell) is thus
calculated by subtracting the E0 for
the oxidation from the E0 for the
reduction.
OR
The half-reaction with the greater potential to be reduced is higher on the table. As
such, the higher half-reaction will happen at the cathode and the lower half-reaction will
occur at the anode (stronger reducing agent therefore more likely to be oxidized).
Chemistry 30 Notes
Zn2+ + 2e-  Zn
Cu2+ + 2e-  Cu
For example: On the table:
E = -0.76V
E = +0.34V
The less positive, or more negative reduction potential, occurs at the anode.
so…Zn + Cu2+  Zn2+ + Cu
E0cell = E0cathode - E0anode
= +0.34V – (-0.76V)
= + 1.10 V
The + sign of the cell potential tells us the redox reaction is spontaneous, meaning the
cell does work.
4. Representing Cell Reactions