Galvanic (or voltaic) cells:
A voltaic cell consists of two half cells, which are connected electrically. A half cell refers to the part of the cell in which one half reaction occurs, and in all text books is depicted as a metal electrode sitting in a beaker of electrolyte solution.
All voltaic cells share several features, which are essential for operation:
The half cell must contain an electrode, (if the reaction involves reduction to or oxidisation of the metal, the electrode will generally be that metal. If neither species in a half reaction is metal, an inert electrode such as platinum will be used). The solution must contain the relevant ions for the redox reaction to occur.
The half cells need to be connected via a wire, so electrons can be transferred from one half cell to the other. The circuit may be connected via a voltmeter, so the potential of the cell can be measured.
If a voltaic cell runs without a salt bridge, charge will quickly build up within the half cells and the cell would stop working. A salt bridge allows ions to flow into each half cell, balancing the charge and allowing extended operation of the cell, but preventing mixing of the two solutions.
Drawing galvanic cells is fairly straightforward – the skeleton two-beaker sketch is a good starting point. We’ll go over how to draw one using an example:
Zn(s) + Cu2+(aq) → Cu(s) + Zn2+(aq)
The first thing to do is break it up into the half reactions and identify which is oxidation and which is reduction:
Zn(s) → Zn2+(aq) + 2e– oxidation
Cu2+(aq) + 2e– → Cu(s) reduction
The anode is the half cell in which oxidation occurs, and reduction occurs in the cathode (an ox and a red cat). The anode is always depicted as being on the left, the cathode the right. Label the anode and the cathode, showing what metals and ions are involved:
Show the direction of electron flow (always anode to cathode) and ions in the salt bridge:
This is a fairly clumsy way of depicting cells, and there is a shortcut known as the “line notation”.
As with sketching a cell, the anode is always on the left and the two half cells are separated by a double line, eg:
Single lines indicate a change in state. The electrodes are always shown on the outer edges, and the reactants are shown in the same order as they would be written in the half reaction. If the half cell includes a gaseous, or two aqueous reactants, an inert electrode is used and any reactants in the same state are separated by a comma, eg:
The cell potential tells you how much of a driving force there is behind the cell reaction, and is the difference between the two half-cell potentials.
Cell potential = oxidation potential + reduction potential
Or… because your potentials are usually given as reduction potentials:
Cell potential = reduction potential of cathode – reduction potential of anode
A note about potentials:
The more positive the standard potential is, the better the species is at oxidising. The more negative, the better the reductant.
The cell potential is easy to work out if you are working in standard conditions, but will change if the concentrations of the components change. Luckily for you, Nernst worked it all out for you and you just need to remember his equation:
Keeping in mind that Q is the reaction quotient and essentially the equilibrium constant (concentration of products/concentration of reactants) when the system isn’t at equilibrium. The reaction is the same for the equilibrium constant, except that at equilibrium the Ecell will be zero.