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Electrolysis - Wikipedia, the fr
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Electrolysis
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This article is about the chemical process. For the cosmetic hair removal procedure, seeElectrology.
Illustration of an electrolysis apparatus used in a school laboratory.
Inchemistryandmanufacturing,electrolysisis a method of using adirect electric current(DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially highly important as a stage in theseparationofelementsfrom naturally occurring sources such asoresusing anelectrolytic cell.
Contents
1 History
2 Overview
2.1 Process of electrolysis
2.2 Oxidation and reduction at the electrodes
2.3 Energy changes during electrolysis
2.4 Related techniques
3 Faraday's laws of electrolysis
3.1 First law of electrolysis
3.2 Second law of electrolysis
4 Industrial uses
5 Competing half-reactions in solution electrolysis
6 Electrolysis of water
7 Electrocrystallization
8 Experimenters
9 See also
10 References
[edit] History
The word electrolysis comes from theGreek λεκτρον [ lektron]"amber" andλ σι [l sis]"dissolution".
1785 Martinus van Marum's electrostatic generator was used to reducetin,zinc, andantimonyfrom their salts using electrolysis.[1]
1800 William NicholsonandJohann Ritterdecomposedwaterintohydrogenandoxygen.
1807 Potassium,sodium,barium,calciumandmagnesiumwere discovered bySir Humphry Davyusing electrolysis.
1875 Paul mile Lecoq de Boisbaudrandiscoveredgalliumusing electrolysis.[2]
1886 Fluorinewas discovered byHenri Moissanusing electrolysis.
1886 Hall-H roult processdeveloped for makingaluminium
1890 Castner-Kellner processdeveloped for makingsodium hydroxide
[edit] Overview
Electrolysis is the passage of adirect electric currentthrough anionicsubstance that is either molten or dissolved in a suitable solvent, resulting in chemical reactions at the electrodes and separation of materials. The main components required to achieve electrolysis are :
Anelectrolyte : asubstancecontaining freeionswhich are the carriers ofelectric currentin theelectrolyte. If theionsare not mobile, as in asolid saltthen electrolysis cannot occur.
Adirect current(DC) supply : provides theenergynecessary to create or discharge theionsin theelectrolyte. Electric current is carried byelectronsin the external circuit.
Twoelectrodes : anelectrical conductorwhich provides the physical interface between theelectrical circuitproviding theenergyand theelectrolyte
Electrodes ofmetal,graphiteandsemiconductormaterial are widely used. Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte and the cost of manufacture.
[edit] Process of electrolysis
The key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons from the external circuit. The required products of electrolysis are in some different physical state from the electrolyte and can be removed by some physical processes. For example, in the electrolysis ofbrineto produce hydrogen and chlorine, the products are gaseous. These gaseous products bubble from the electrolyte and are collected.[3]
2 NaCl + 2 H2O → 2 NaOH + H2+ Cl2
A liquid containing mobile ions (electrolyte) is produced by:
Solvationor reaction of anionic compoundwith asolvent(such as water) to produce mobile ions
An ionic compound is melted (fused) by heating
An electrical potential is applied across a pair ofelectrodesimmersed in the electrolyte. Each electrode attracts ions that are of the oppositecharge. Positively charged ions (cations) move towards the electron-providing (negative) cathode, whereas negatively charged ions (anions) move towards the positive anode. At the electrodes,electronsare absorbed or released by the atoms and ions. Those atoms that gain or lose electrons to become charged ions pass into the electrolyte. Those ions that gain or lose electrons to become uncharged atoms separate from the electrolyte. The formation of uncharged atoms from ions is called discharging. The energy required to cause the ions to migrate to the electrodes, and the energy to cause the change in ionic state, is provided by the external source of electrical potential.
[edit] Oxidation and reduction at the electrodes
Oxidationof ions or neutral molecules occurs at theanode, and thereductionof ions or neutral molecules occurs at thecathode. For example, it is possible to oxidize ferrous ions to ferric ions at the anode:
Fe2+
aq→ Fe3+
aq+ e
It is also possible to reduceferricyanideions toferrocyanideions at the cathode:
Fe(CN)3-
6+ e → Fe(CN)4-
6
Neutral molecules can also react at either electrode. For example: p-Benzoquinone can be reduced to hydroquinone at the cathode: + 2 e + 2 H+→ In the last example, H+ions (hydrogen ions) also take part in the reaction, and are provided by an acid in the solution, or the solvent itself (water, methanol etc.). Electrolysis reactions involving H+ions are fairly common in acidic solutions. In alkaline water solutions, reactions involving OH-(hydroxide ions) are common. The substances oxidised or reduced can also be the solvent (usually water) or the electrodes. It is possible to have electrolysis involving gases.
[edit] Energy changes during electrolysis
The amount of electrical energy that must be added equals the change inGibbs free energyof the reaction plus the losses in the system. The losses can (in theory) be arbitrarily close to zero, so the maximumthermodynamicefficiency equals theenthalpychange divided by the free energy change of the reaction. In most cases, the electric input is larger than the enthalpy change of the reaction, so some energy is released in the form of heat. In some cases, for instance, in the electrolysis ofsteaminto hydrogen and oxygen at high temperature, the opposite is true. Heat is absorbed from the surroundings, and theheating valueof the produced hydrogen is higher than the electric input.
[edit] Related techniques
The following techniques are related to electrolysis:
Electrochemical cells, including the hydrogenfuel cell, utilise differences inStandard electrode potentialin order to generate an electrical potential from which useful power can be extracted. Although related via the interaction of ions and electrodes, electrolysis and the operation of electrochemical cells are quite distinct. A chemical cell shouldnotbe thought of as performing "electrolysis in reverse".
[edit] Faraday's laws of electrolysis
Main article:Faraday's laws of electrolysis
[edit] First law of electrolysis
In 1832,Michael Faradayreported that the quantity of elements separated by passing an electric current through a molten or dissolvedsaltis proportional to the quantity of electric charge passed through the circuit. This became the basis of the first law of electrolysis:
[edit] Second law of electrolysis
Faraday discovered that when the same amount of electricity is passed through different electrolytes, the mass of substance liberated/deposited at the electrodes is directly proportional to their equivalent weights.
[edit] Industrial uses
Hall-Heroult processfor producingaluminium
Production ofaluminium,lithium,sodium,potassium,magnesium,calcium
Coulometrictechniques can be used to determine the amount of matter transformed during electrolysis by measuring the amount of electricity required to perform the electrolysis
Production ofchlorineandsodium hydroxide
Production ofsodium chlorateandpotassium chlorate
Production of perfluorinated organic compounds such astrifluoroacetic acid
Production ofelectrolytic copperas acathode, from refinedcopperof lower purity as ananode.
Electrolysis has many other uses:
Electrometallurgyis the process of reduction of metals from metallic compounds to obtain the pure form of metal using electrolysis. For example, sodium hydroxide in its molten form is separated by electrolysis into sodium and oxygen, both of which have important chemical uses. (Water is produced at the same time.)
Anodizationis an electrolytic process that makes the surface of metals resistant tocorrosion. For example, ships are saved from being corroded by oxygen in the water by this process. The process is also used to decorate surfaces.
Abatteryworks by the reverse process to electrolysis.
Production ofoxygenforspacecraftandnuclear submarines.
Electroplatingis used in layering metals to fortify them. Electroplating is used in many industries for functional or decorative purposes, as in vehicle bodies and nickel coins.
Production of hydrogen for fuel, using a cheap source of electrical energy.
Electrolytic Etching of metal surfaces like tools or knives with a permanent mark or logo.
Electrolysis is also used in the cleaning and preservation of old artifacts. Because the process separates the non-metallic particles from the metallic ones, it is very useful for cleaning old coins and even larger objects.
[edit] Competing half-reactions in solution electrolysis
Using theNernst equationtheelectrode potentialcan be calculated for a specific concentration of ions, temperature and the number of electrons involved. For pure water (pH7):
the electrode potential for the reduction producing hydrogen is －0.41 V
the electrode potential for the oxidation producing oxygen is +0.82 V.
Comparable figures calculated in a similar way, for 1Mzinc bromide, ZnBr2, are －0.76 V for the reduction to Zn metal and +1.10 V for the oxidation producing bromine. The conclusion from these figures is that hydrogen should be produced at the cathode and oxygen at the anode from the electrolysis of water which is at variance with the experimental observation that zinc metal is deposited and bromine is produced.[6]The explanation is that these calculated potentials only indicate the thermodynamically preferred reaction. In practice many other factors have to be taken into account such as the kinetics of some of the reaction steps involved. These factors together mean that a higher potential is required for the reduction and oxidation of water than predicted, and these are termedoverpotentials. Experimentally it is known thatoverpotentialsdepend on the design of the cell and the nature of the electrodes. For the electrolysis of a neutral (pH 7) sodium chloride solution, the redu
ction
of sodium ion is thermodynamically very difficult and water is reduced evolving hydrogen leaving hydroxide ions in solution. At the anode the oxidation of chlorine is observed rather than the oxidation of water since the overpotential for the oxidation ofchloridetochlorineis lower than the overpotential for the oxidation ofwatertooxygen. Thehydroxide ionsand dissolvedchlorinegas react further to formhypochlorous acid. The aqueous solutions resulting from this process is calledelectrolyzed waterand is used as a disinfectant and cleaning agent.
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