R3 The classic electrochemical cell

In illustration R2, it was shown that the electron transfer between I^- and iron(III) ions (Fe³⁺) can be used to produce an electric current by placing platinum electrodes in the two solutions, connecting the electrodes with a wire and connecting the solutions with a salt bridge. The two half-reactions, which make up the electron transfer reaction, take place in the solutions in the two beakers.
This is an electrochemical cell, and the beakers in which the two half-reactions take place are its half-cells. An electrochemical cell is therefore a device which allows an electron transfer reaction to take place in such a way that a flow of electrons is produced between the electrodes.

If a piece of zinc metal is placed in a blue aqueous solution of copper(II) sulphate (CuSO₄) then the colour intensity spontaneously decreases and the zinc takes on a red-brown colour. An electron transfer takes place between the zinc metal and the copper(II) ions with copper being deposited on the zinc metal, the solution becoming colourless as the copper(II) ions are reduced on the zinc metal to copper and the zinc metal is oxidised to colourless zinc ions. This reaction can be represented as:

This electron transfer can, in principle, be observed indirectly by an increase in temperature as the electron transfer energy is changed to heat energy.

This redox reaction can be broken down into the following half-reaction equations :

If these half-reactions are set-up in separate beakers as in illustration R2 (right-hand side), each containing the metals in the form of electrodes immersed in an aqueous solution of the corresponding metal salt and with the solution connected by a salt bridge and the metal electrodes linked by wires to a lamp as shown in illustration R3, we again create a source of electric current.

This demonstrates the involvement of electron transfer.
If this simple electrochemical cell is kept at a temperature of 25°C (298K) and the concentrations of the zinc and copper ions are both 1mol/L, then we obtain what is known as the Daniell cell.

The tendency of electrons to flow through the external circuit of a cell is quantified as the e.m.f. , the electromotive force of a cell, also called the cell potential E_cell.

The e.m.f of a cell can be measured using a high resistance voltmeter, which takes negligible current from the cell.
In the case of the Daniell cell the E_cell is +1.1 V.
The cell potential is by convention taken to be acting from left to right through the cell with the (more) negative electrode on the left.
Thus

The cell can be represented as :

In practice sulphate salts are used in this cell although nitrate, chloride or hydrogen sulphate salts could be used.

The salt bridge, as also shown in illustration R2, contains KCl or KNO₃. It is a sufficiently good conductor of ions to produce good contact between the solutions.
It does not contain possibly interfering species.

When the electrodes are connected, current begins to flow and the reaction starts. As the electrons flow between the zinc and copper electrodes the concentration of zinc ions increases in one half-cell while the concentration of copper(II) ions decreases in the other half-cell leading to an ionic imbalance.

This results in a decrease in the potential difference between the two half-cells. The salt bridge between the two solutions enables an ionic balance to be maintained without the transport of ions from one half-cell solution to the other. The positive ions (cations) of the salt bridge diffuse to the sulphate solution and the negative ions (anions) to the solution of zinc ions.
The reaction proceeds and the potential difference between the electrodes decreases eventually reaching zero when either all the copper(II) ions have or the zinc electrode has been totally consumed.

On the right-hand side of illustration R3 a lemon is shown with two electrodes stuck in it. One is zinc and one is copper. If these electrodes are joined by a wire to a voltmeter then a voltage (admittedly lower than 1.1 V) can be measured. Again the zinc is functioning as the anode and the copper as the cathode.

The lemon must contain ions which function in the same way as the Cu²⁺ and Zn²⁺ ions in the Daniell cell. The diffusion of the ions is slow and so the voltage is lower but is measurable, even though both electrodes are in one solution without a salt bridge.

Cell electrodes

A metal consists of cations of the metal (positively charged ions) surrounded by electrons. If a strip of metal is placed in a solution of its ions some of the cations in the metal may dissolve leaving a build up of electrons on the metal:

The metal strip will become negatively charged.

Alternatively metal ions in the solution may take electrons from the strip of metal and be discharged as metal atoms:

The potential difference between the strip of metal and the solution depends on the nature of the metal and the concentration of the ions involved in the equilibrium at the metal’s surface. If we compare zinc (Zn) and copper (Cu) for example, then, for the same concentration of ions in each solution, zinc acquires a more negative potential than copper since it has a greater tendency to dissolve ions and a smaller tendency to be deposited as a metal. Hence in such a system as is shown in illustration R3 the zinc electrode is the anode and the copper electrode the cathode.
At the anode oxidation takes place : here zinc metal is converted into zinc ions.
At the cathode reduction takes place : copper(II) ions are reduced to copper metal.
In galvanic cells, the anode is indicated by the negative sign s and the cathode by the sign r.
In electrolytic cells (see R11) an anode, being the electrode where oxidation takes place, gets the positive sign r. A cathode of an electrolytic cell is indicated by the negative sign s.
For a galvanic cell, the cell voltage or cell potential, in terms of anode and cathode is written as :