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Electrochemistry is the study of interchange of chemical and electrical energy.

Oxidation/Reduction involves the exchange of electrons from one chemical species to another.

Normally, this is done when the two chemicals contact each other in the activated complex (when two

species bump into each other in solution for example).

We are interested in separating the chemical species such that the electrons transfer via an external

circuit. That way, we can measure the electrochemical effects.

To properly understand the connection between the redox reaction and the electricity, we should

balance the overall redox reaction using a half-reaction method such as the one described in the

previous section of these notes. We can set up the physical reaction vessel such that the chemicals from

one half reaction are separated from those of the second half reaction. For reaction to occur, we still

need to connect the solutions to complete the circuit. This is done by attaching wires between

electrodes in the two half cells and by connecting the solutions of the two half cells via a salt bridge or

by some other device such as a semi-permeable membrane.

In general, such a cell is called an electrochemical cell. These cells could be used in one of two types of

situations:

1. The chemical reaction is spontaneous and produces electricity.

This is called a voltaic cell or a galvanic cell.

2. The chemical reaction is non-spontaneous and is forced by electricity from an external source.

This kind of cell is called an electrolysis cellGalvanic or Voltaic cells

Consider a piece of zinc foil placed in a beaker of copper sulphate. The copper sulphate solution is blue

because of the presence of the Cu2+ ions. When the zinc is added, the solution changes to colourless

and the zinc metal is dissolved and replaced by a reddish orange powder. The colourless solution no

longer has Cu2+ in it and the reddish orange powder is Cu(s). The Zn(s) is dissolved and is now Zn2+. [I

know this because of my vast knowledge of chemistry ;-)]

If we write an overall reaction for this process, we get:

Zn + Cu2+ Zn2+ + Cu. This doesn't help us in our quest for electrochemistry knowledge. Let's rewrite

this in half-reaction form.

Zn Zn2+ + 2e

Cu2+ + 2e Cu

Now, we can set up two half cells, one with a zinc electrode in a Zn2+ solution (say ZnSO4) and the other

with a copper electrode in a Cu2+ solution (say CuSO4) as follows.

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The centre line in the diagram, recall, is either a semi-permeable membrane or a salt bridge.

Now, at the anode, we have the reaction

Zn(s) Zn2+(aq) + 2e(aq)

and at the cathode, we have

Cu2+(aq) + 2e(aq) Cu(s)

Electrons pass through the wires and SO42 pass through the membrane to keep the solutions

neutralCell Potentials

Electrons from solution pushed onto the anode, around the external circuit and onto the cathode where

they are pulled out into the solution.

That's one way of thinking of the electrical circuit part of the electrochemical cell. The electrons pushed

around the external circuit can do work (run a motor, illuminate a light bulb, etc). The amount of work

possible is a function of both the voltage (potential) and of the current (number of electrons) in the

circuit.

Pushing one Coulomb of charge around a circuit at a potential of 1 Volt does One Joule of work. OR,

mathematically, 1J = 1V 1C.

The cell potential (voltage of the cell) depends on the chemicals used. For example, the chemicals in dry-

cells (batteries) are such that the potential is always about 1.5 V. This has become a standard and is now

a limiting factor in deciding which chemicals can be used to create a battery.

The cell potential is given a symbol of Ecell. If all chemicals are at activity of 1 (conc. = 1 M, p = 1 bar)

then the cell potential is the standard cell potential and is given as Ecell.

Any redox reaction has the potential (pun) to be used in an electrochemical cell. We merely need to be

able to divide the oxidizing and reducing agents into two half cells (half reactions).

Take for example, the reaction of zinc metal dissolving in hydrochloric acid. The reaction is:

Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)

We need to separate the zinc from the hydrogen. We can use zinc as an electrode but what about the

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hydrogen.

In this case, we need to set up a special electrode, which allows H2 gas molecules to interact directly

with H+ dissolved in water and with the electrons from the external circuit simultaneously. Such a

system is pictured below.

A blow-up of the surface of the platinum electrode is shown below so the location of reaction can be

better understood.

The other half-cell would look much like that pictured in a previous diagram. The whole cell diagram, of

course would include the external part of the circuit and the salt bridge or membrane to complete the

circuit. We can abbreviate this diagram as follows:

Zn(s)/ZnSO4(aq)//H2SO4(aq)/H2(g),Pt(s)

[Anode // Cathode]

Where the single slash mark / represents the boundary between solution and electrode and the double

slash // represents the salt bridge or semi-permeable membrane. The external circuit, of course, joins

the two electrodes (solid) and is not explicitly shown here.

The overall cell voltage can be summed from the half-cell potentials of the oxidation and of the

reduction reactions.

Ecell = ERed + Eox