Chemistry » Chemistry 111: Electrochemical Reactions » Applications Of Electrochemistry

Summary and Main Ideas

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Summary

  • Oxidation is the loss of electrons and reduction is the gain of electrons.

  • A redox reaction is one where there is always a change in the oxidation numbers of the elements that are involved in the reaction.

  • It is possible to balance redox equations using the half-reactions that take place within the overall reaction.

  • An electrochemical reaction is one where either a chemical reaction produces an electric current, or where an electric current causes a chemical reaction to take place.

  • In a galvanic cell a chemical reaction produces a current in the external circuit. An example is the zinc-copper cell.

  • An electrolytic cell is an electrochemical cell that uses electricity to drive a non-spontaneous reaction. In an electrolytic cell, electrolysis occurs, which is a process of separating elements and compounds using an electric current.

  • Cells have a number of components. They consist of two electrodes, which are connected to each other by an external circuit wire.

  • In a galvanic cell each electrode is placed in a separate container in an electrolyte solution. The two electrolytes are connected by a salt bridge.

  • In an electrolytic cell both electrodes are placed in the same container in an electrolyte solution.

  • One of the electrodes is the anode, where oxidation takes place. The cathode is the electrode where reduction takes place.

  • In a galvanic cell, the build up of electrons at the anode sets up a potential difference between the two electrodes, and this causes a current to flow in the external circuit.

  • Standard cell notation for a galvanic cell has the anode on the left and the cathode on the right. For example:

    \(\text{Zn}(\text{s})|\text{Zn}^{2+}(\text{aq})||\text{Cu}^{2+}(\text{aq})|\text{Cu}(\text{s})\)

    \(|\) = a phase boundary (solid/aqueous)

    \(||\) = the salt bridge

  • Different metals have different reaction potentials. The reduction potential of metals (in other words, their ability to ionise), is recorded in a table of standard electrode reduction potentials. The more negative the value, the greater the tendency of the metal to be oxidised. The more positive the value, the greater the tendency of the metal to be reduced.

  • The values on the table of standard electrode potentials are measured relative to the standard hydrogen electrode.

  • The EMF of a cell can be calculated using one of the following equations:

    E°(cell) = E°(reduction half-reaction) – E°(oxidation half-reaction)

    E°(cell) = E°(oxidising agent) – E°(reducing agent)

    E°(cell) = E°(cathode) – E°(anode)

  • It is possible to predict whether a reaction is spontaneous or not, either by looking at the sign of the cell EMF or by comparing the electrode potentials of the two half-cells.

  • Industrial applications of cells include electrolysis (the electrowinning of copper), in the chloralkali industry (mercury, diaphragm and membrane cells), as well as the extraction of metals from ores (e.g. aluminium from bauxite).

A comparison of galvanic and electrolytic cells

It should be much clearer now that there are a number of differences between galvanic and electrolytic cells. Some of these differences have been summarised in the table below.

 

Galvanic cell

Electrolytic cell

Chemical

reactions

spontaneous reactions

non-spontaneous reactions

Energy

changes

Chemical potential

energy from chemical

reactions is converted to

electrical energy

An external supply of

electrical energy causes a

chemical reaction to occur

Anode

is negative, oxidation

occurs at anode

is positive, oxidation

occurs at anode

Cathode

is positive, reduction

occurs at cathode

is negative, reduction

occurs at cathode

Cell set-up

two half-cells, one

electrode in each,

connected by a salt-bridge

one cell, both electrodes in

cell, no salt-bridge

Electrolyte

solution(s)

The electrolyte solutions

are kept separate from

one another, and are

connected by a salt bridge

The cathode and anode are

in the same electrolyte

Applications

batteries

Electrolysis e.g. of water,

NaCl, electroplating

Table: A comparison of galvanic and electrolytic cells.

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