Electrolysis

23.4K views03:00 min

Spontaneous redox reactions in voltaic cells convert chemical energy into electrical energy. Alternatively, external electrical energy drives nonspontaneous redox reactions in a process called electrolysis, taking place in an electrolytic cell.

An electrolytic cell has an anode for oxidation and a cathode for reduction. However, unlike the voltaic cell, where the negatively charged anode is the source of electrons, the electrolytic cell has an external source of electrons like a battery.

The battery draws electrons from the anode connected to its positive terminal and drives electrons towards the cathode, which is connected to its negative terminal. Thus electrons flow from the positive to the negative terminal of the external power source.

Electrolysis is used to separate ionic compounds into their constituent elements. In their molten or aqueous state, the constituent ions are free to conduct electricity.

Consider potassium bromide that melts at a temperature of 734 °C. Here, bromide and potassium ions are present at their lowest or highest oxidation state, respectively. Under power, the bromide ions are oxidized to bromine gas at the anode, and potassium ions are reduced to solid potassium at the cathode.

In molten ionic mixtures like sodium iodide and potassium bromide, standard electrode potentials are used to predict which substance will react. Generally, anions with a lower electrode potential are oxidized, and cations with a more positive electrode potential are reduced.

Interestingly, electrolysis of aqueous potassium bromide produces different products from molten salts due to the possible electrolysis of water.

Thus, two reactions could occur at each electrode: the oxidation of bromide ions or water at the anode and the reduction of potassium ions or water at the cathode.

Given the lower electrode potential, oxidation of bromide ions to bromine is favored, while the more positive electrode potential of water drives the reduction to hydrogen gas.

Stoichiometry of the half-reactions can be used to calculate the amount of substances produced or consumed during electrolysis.

For example, electrolysis of molten potassium bromide uses one electron to form a potassium atom. Thus, every mole of electrons, or 96,485 coulombs of charge, passed through the cell produces one mole or nearly 39 grams of solid potassium.

Overview

$$\rightleftharpoonup{xx}$$ $$\longleftharp{xx}$$, $$\longrightharp{xx}$$,

In a galvanic cell, the electrical work is done by a redox system on its surroundings as electrons produced by the spontaneous redox reactions are transferred through an external circuit. Alternatively, an external circuit does work on a redox system by imposing a voltage sufficient to drive an otherwise nonspontaneous reaction in a process known as electrolysis. For instance, recharging a battery involves the use of an external power source to drive the spontaneous (discharge) cell reaction in the reverse direction, restoring to some extent the composition of the half-cells and the voltage of the battery. Other examples include the use of electrolysis in the refinement of metallic ores, the manufacture of commodity chemicals, and the electroplating of metallic coatings onto utensils, jewelry, etc.

Predicting the Product of Electrolysis

The electrolysis of molten sodium chloride, NaCl (l), is used for the industrial production of metallic sodium, Na, and chlorine gas, Cl₂. Sodium ions (Na⁺) are reduced to atoms at the cathode, while chloride ions (Cl⁻) ions are oxidized to chlorine gas, Cl₂, at the anode. The redox reactions are:

$Electric dipole configuration; diagram with charge $q$ separation $d$ and dipole moment direction.$

The negative cell potential indicates a nonspontaneous reaction that must be driven by imposing a positive potential of magnitude greater than −4.07 V. Electrolysis of molten sodium chloride is carried out at a high temperature of 801 °C as ionic solids melt at high temperatures.

The electrolysis of water produces stoichiometric amounts of oxygen gas at the anode and hydrogen at the cathode. To improve electrical conductivity, the hydrogen ion concentration of the water is increased by adding a strong acid. The associated redox reactions are:

Chemical equilibrium; reaction equation CH4 + 2O2 = CO2 + 2H2O; combustion process diagram.

The electrolysis of solutions of ionic compounds such as aqueous sodium chloride may involve the electrolysis of either water species (H₂O, H⁺, OH^-) or solute species (the cations and anions of the compound) at the anode and cathode.

The electrolysis of aqueous sodium chloride could involve either of these two anode reactions:

Chemical equilibrium equations, ΣMA=0, formula diagram, statics education, reaction dynamics.

Thermodynamically, water would be more readily oxidized; however, in practice, chlorine gas is produced. Oxidation of water would require a much larger voltage to initiate. To overcome this overvoltage, electrodes are chosen and the potential of the cell is carefully monitored to ensure that the oxidation of chloride ions at the anode.

Similarly, the possible reduction reactions at the cathode are:

DNA replication process diagram with leading and lagging strands, highlighting Okazaki fragments.

Comparing these standard half-reaction potentials suggests the reduction of water is thermodynamically favored. The net cell reaction in this case is then:

Photosynthesis, light-dependent reactions diagram; chlorophyll, ATP synthesis, electron transport chain.

Quantitative Aspects of Electrolysis

Electrical current is the rate of flow of electrons and is measured in ampere, one coulomb per second (A = 1 C/s). The charge transferred, Q, by passage of a constant current, I, over a specified time interval, t, is given by

Equilibrium equations diagram, ΣFx=0, ΣFy=0, static balance analysis, physics study.

When electrons are transferred during a redox process, the stoichiometry of the reaction may be used to derive the total amount of (electronic) charge involved. For example, the generic reduction process,

Equilibrium diagram with ΣFx=0, ΣFy=0, depicting static equilibrium conditions in mechanics.

involves the transfer of n mole of electrons. The charge transferred is, therefore,

Static equilibrium diagram, hydrostatic system, ΣFy=0, forces labeled, physics concept.

where F is Faraday’s constant, the charge in coulombs for one mole of electrons. For an electrochemical cell, the current flow is measured and can be used in stoichiometric calculations related to the cell reaction.