Voltaic (Galvanic) cells
● Voltaic cells convert energy from spontaneous, exothermic chemical processes to electrical energy.
● Oxidation occurs at the anode (negative electrode) and reduction occurs at the cathode (positive
electrode) in a voltaic cell.
● A voltaic cell generates an electromotive force (EMF) resulting in the movement of electrons from
the anode (negative electrode) to the cathode (positive electrode) via the external circuit. The EMF
is termed the cell potential (E *).
● The standard hydrogen electrode (SHE) consists of an inert platinum electrode in contact with
1 mol dm–3 hydrogen ion and hydrogen gas at 100 kPa and 298 K. The standard electrode potential
(E *) is the potential (voltage) of the reduction half-equation under standard conditions measured
relative to the SHE. Solute concentration is 1 mol dm–3 or 100 kPa for gases. E * of the SHE is 0 V.
● G * = –nFE *. When E * is positive, G * is negative indicative of a spontaneous process. When E * is
negative, G * is positive indicative of a non-spontaneous process. When E * is 0, then G * 0.
Guidance
● G * = –nFE * is given in the data booklet in section 1.
● Faraday’s constant = 96 500 C mol−1 is given in the data booklet in section 2.
Electrolytic cells
● Electrolytic cells convert electrical energy to chemical energy, by bringing about non-spontaneous
processes.
● Oxidation occurs at the anode (positive electrode) and reduction occurs at the cathode (negative
electrode) in an electrolytic cell.
● When aqueous solutions are electrolysed, water can be oxidized to oxygen at the anode and
reduced to hydrogen at the cathode.
● Current, duration of electrolysis, and charge on the ion affect the amount of product formed at the
electrodes during electrolysis.
Guidance
Electrolytic processes to be covered in theory should include the electrolysis of aqueous solutions (e.g.
sodium chloride, copper(II) sulfate, etc.) and water using both inert platinum or graphite electrodes and
copper electrodes.
● Electroplating involves the electrolytic coating of an object with a metallic thin layer.
Guidance
The term cells in series should be understood.

● Construction and annotation of both types of electrochemical cells.
Guidance
For voltaic cells, a cell diagram convention should be covered.
● Explanation of how a redox reaction is used to produce electricity in a voltaic cell and how current
is conducted in an electrolytic cell.
● Distinction between electron and ion fl ow in both electrochemical cells.
● Performance of laboratory experiments involving a typical voltaic cell using two metal/metal-ion
half-cells.
● Deduction of the products of the electrolysis of a molten salt.
Guidance
Explanations should refer to E * values, nature of the electrode, and concentration of the electrolyte.
● Calculation of cell potentials using standard electrode potentials.
● Prediction of whether a reaction is spontaneous or not using E * values.
● Determination of standard free-energy changes (G *) using standard electrode potentials.
● Explanation of the products formed during the electrolysis of aqueous solutions.
● Perform lab experiments that could include single replacement reactions in aqueous solutions.
● Determination of the relative amounts of products formed during electrolytic processes.
● Explanation of the process of electroplating.

Is there a difference between an arbitrary decision and a random decision? Would it make a difference to the study of electrochemistry if a different half-cell, for example Zn2+(aq)/Zn(s), was used as the reference standard for electrode potentials?

Task: 
1. Watch the videos for the introduction to electrochemistry, voltaic cells and electrolytic cells activities below on this website
2. As you learn more about electrochemistry add to the overview