Electrochemistry is a branch of chemistry that describes and measures the relationship between electrical energy and a chemical change. Electrochemical reactions involve the movement of electrons from one species to another. If the reaction is spontaneous, it can result in a generated current. If the reaction is not spontaneous, it can be driven by the application of current.
In electrochemistry, the key reaction is the oxidation-reduction reaction, called a redox reaction. The redox reaction is composed of two half-reactions; oxidation, where a substance known as the reducing agent loses electrons, and reduction, where a substance known as the oxidizing agent gains electrons. Redox reactions always occur in pairs and change the oxidation states of the atoms in the molecules involved. An easy way to remember this relationship is through the pneumatic OIL-RIG: Oxidation Is Losing – Reduction Is Gaining.
An electrochemical cell is used to measure or induce electrochemical reactions. It consists of several components: a chamber containing the reaction solution, two conductive electrodes, a conductive electrolyte solution, and an external circuit. There are two types of electrochemical cells. One is the electrolytic cell, which uses electrical energy to drive a nonspontaneous reaction. In this type of cell, electrical energy is supplied from an external power source.
The other type of cell is the galvanic cell, which uses a spontaneous electrochemical reaction to generate electrical energy. The two electrodes are called the anode and cathode, and they are connected by an external circuit. The reaction chamber is filled with an electrolyte, which promotes the passage of ions between electrodes. In a galvanic cell, a salt bridge completes the circuit by allowing the flow of ions between the electrolytic solutions containing the electrodes. In an electrolytic cell, there is no salt bridge as the electrodes are typically in the same electrolytic solution.
The reduction reaction occurs at the cathode, while the oxidation reaction occurs at the anode. This is remembered using the mnemonic “Red Cat,” which means that reduction occurs at the cathode.
Electrolytic cells typically have one reaction chamber, which contains the electrolyte solution. The electrolyte is usually an aqueous solution containing ions or dissolved salts. The ions in the electrolyte promote the movement of ions and electrons through the solution.
When an external voltage is applied, ions in the electrolyte are attracted to the electrode with the opposite charge. This is where the two half-reactions occur. The anode loses electrons during oxidation, while the cathode gains electrons during reduction.
Electroplating is a process that utilizes an electrolytic cell. Electroplating is a process in which one metal is deposited on the surface of an electrode, typically another metal. British scientist Michael Faraday, for whom the Faraday constant is named after, demonstrated the molar relationship between the electroplating charged ion and the electrical current. Looking at the specific half-reaction for silver:
Ag+ + 1e- → Ag
It takes one mole of electrons, supplied by an external current, to reduce one mole of silver cations into solid silver. Therefore, half-reaction stoichiometry can determine the amount of the electroplating material utilized for a known quantity of electrons. Recall that electricity is not measured based on the moles of electrons but rather in coulombs, named after the French engineer Charles-Augustin de Coulomb. One coulomb is equal to 1 amp per second, which is based on the principle of elementary charge, or the charge a single electron carries. Faraday’s discovery can therefore relate electrical current to the number of moles the electroplating cation plated:
F = eNA
F is the Faraday constant, e is the elementary charge expressed in coulombs, and NA is Avogadro’s number. Therefore, Faraday’s constant is expressed as the number of coulombs per mole and has a value of 96485 coulombs per mole. How can we determine the number of moles of electrons transferred to the anode based on the current? Using the electrical charge Q and the definition of a coulomb:
The electrical charge is equal to the current in amps (I) multiplied by the time in seconds that the current was allowed to flow. By dividing Q by the Faraday constant, which has the units of coulombs per mole:
This equation allows us to calculate the number of moles of electrons and thus determine how much of the electroplating cation was reduced.
In electroplating, the metal comprises the anode plates, or covers, the cathode in a thin layer of metal. The amount of metal plated is dependent on the amount of current applied as well as the number of moles of the electroplating cation available. An external power source, such as a battery, induces the flow of electrons from the anode to the cathode, and from the positive terminal to the negative terminal on the battery.
For example, consider an electrolytic cell with a copper electrode, a brass key acting as the second electrode, and an aqueous solution of copper sulfate as the electrolyte. Here, copper from the electrolyte and the copper electrode is plated onto the brass key.
For the copper metal to be deposited on the brass key, the solid copper electrode must be oxidized to form copper ions. Then, the copper cations from both the electrode and electrolyte are reduced from the solution to form solid copper on the brass key.
Cu2+ + 2e- → Cu(s)
The electrons for the reaction are received from the negative terminal of the battery. Thus, the reduction reaction occurs at the brass key, while the oxidation reaction occurs at the copper electrode. The concentrated and acidified copper sulfate solution increases solubility; therefore, the higher the concentration of the solution, the lower the resistance and the higher the current. In turn, a higher current allows more plating of the copper ions onto the brass key to occur.
Some metals have a greater tendency than others to lose electrons. The standard electrode potential (E°) of a substance is the measure of the tendency of a substance to lose electrons. The metal with the highest reduction potential has the highest tendency to lose electrons; thus, it is electroplated first.
- Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Chemistry and Chemical Reactivity. Belmont, CA: Brooks/Cole, Cengage Learning.
- Silberberg, M.S. (2009). Chemistry: The Molecular Nature of Matter and Change. Boston, MA: McGraw Hill.