The study of chemical change in a solution, resulting from the uptake of electrical energy from an external circuit or its supply to a circuit. Storage batteries (accumulators) illustrate this process, as they charge and discharge respectively. In all cases, changes in the oxidation states of elements occur, and energy is converted between chemical and electrical forms.
Electrochemistry is a branch of chemistry that studies the reactions which take place at the interface of an electronic conductor (the electrode composed of a metal or a semiconductor, including graphite) and an ionic conductor (the electrolyte).
If a chemical reaction is caused by an external voltage, or if a voltage is caused by a chemical reaction, as in a battery, it is an electrochemical reaction.
History
16th to 18th century developments
The 16th century marked the beginning of the electrical understanding.
In 1663 the German physicist Otto von Guericke created the first electric generator, which produced static electricity by applying friction in the machine.
By the mid—1700s the French chemist Charles François de Cisternay du Fay discovered two types of static electricity, and that like charges repel each other whilst unlike charges attract.
Charles-Augustin de Coulomb developed the law of electrostatic attraction in 1781 as an outgrowth of his attempt to investigate the law of electrical repulsions as stated by Joseph Priestley in England.
In the late 1700s the Italian physician and anatomist Luigi Galvani marked the birth of electrochemistry by establishing a bridge between chemical reactions and electricity on his essay "De Viribus Electricitatis in Motu Musculari Commentarius" (Latin for Commentary on the Effect of Electricity on Muscular Motion) in 1791 where he proposed a "nerveo-electrical substance" on biological life forms.
On his essay Galvani concluded that animal tissue contained a here-to-fore neglected innate, vital force, which he termed "animal electricity," which activated nerves and muscles spanned by metal probes.
Galvani's scientific colleagues generally accepted his views, but Alessandro Volta rejected the idea of an "animal electric fluid," replying that the frog's legs responded to differences in metal temper, composition, and bulk.
19th century
In 1800, the English chemists William Nicholson and Johann Ritter succeeded in decomposing water into hydrogen and oxygen by electrolysis.
Hans Christian Ørsted's discovery of the magnetic effect of electrical currents in 1820 was immediately recognized as an epoch-making advance, although he left further work on electromagnetism to others.
In 1821, Estonian-German physicist Thomas Johann Seebeck demonstrated the electrical potential in the juncture points of two dissimilar metals when there is a heat difference between the joints.
In 1827 the German scientist Georg Ohm expressed his law in this famous book "Die galvanische Kette, mathematisch bearbeitet" (The Galvanic Circuit Investigated Mathematically) in which he gave his complete theory of electricity.
In 1832 Michael Faraday's experiments on Electrochemistry led him to state his two laws of electrochemistry.
William Grove produced the first fuel cell in 1839.
In 1894 Friedrich Ostwald concluded important studies of the electrical conductivity and electrolytic dissociation of organic acids.
Hermann Nernst developed the theory of the electromotive force of the voltaic cell in 1888.
In 1898 Fritz Haber showed that definite reduction products can result from electrolytic processes if the potential at the cathode is kept constant.
The 20th century and recent developments
In 1909, Robert Andrews Millikan began a series of experiments to determine the electric charge carried by a single electron.
In 1923, Johannes Nicolaus Brønsted and Thomas Martin Lowry published essentially the same theory about how acids and bases behave, using an electrochemical basis.
Principles
Redox reactions
Electrochemical process are redox reactions where energy is produced by a spontaneous reaction which produces electricity, otherwise electrical current stimulates a chemical reaction.
Oxidation and Reduction
The elements involved in an electrochemical reaction are characterized by the number of electrons each has.
For example when sodium reacts with chlorine, sodium donates one electron and gains an oxidation state of +1.
The loss of electrons of a substance is called oxidation, and the gain of electrons is reduction.
The substance which loses electrons is also known as the reducing agent, or reductant, and the substance which accepts the electrons is called the oxidizing agent, or oxidant.
The gain of oxygen, loss of hydrogen and increase in oxidation number is also considered to be oxidation, while the inverse is true for reduction.
A reaction in which both oxidation and reduction is occurring is called a redox reaction.
Oxidation requires an oxidant.
Balancing redox reactions
Electrochemical reactions in water are better understood by balancing redox reactions using the Ion-Electron Method where H ion, H2O and electrons (to compensate the oxidation changes) are added to cell's half reactions for oxidation and reduction.
Acid medium
In acid medium H atoms and water are added to half reactions to balance the overall reaction.
Finally the reaction is balanced by multiplying the number of electrons from the reduction half reaction to oxidation half reaction and vice versa and adding both half reactions, thus solving the equation.
Reaction balanced:
Basic medium
In basic medium OH- ions and water are added to half reactions to balance the overall reaction.
The same procedure as followed on acid medium by multiplying electrons to opposite half reactions solve the equation thus balancing the overall reaction.
Equation balanced:
Neutral medium
The same procedure as used on acid medium is applied, for example on balancing using electron ion method to complete combustion of propane gas.
As in acid and basic medium, electrons which were used to compensate oxidation changes are multiplied to opposite half reactions, thus solving the equation.
Equation balanced:
Electrochemical cells
An electrochemical cell is a device capable of producing electric current from energy released by a spontaneous redox reaction.
In a Galvanic cell the anode is defined as the electrode where oxidation occurs and the cathode is the electrode where the reduction takes place.
The Galvanic cell's metals dissolve in the electrolyte at two different rates, leaving some electrons in the rest of the metal, which makes it negative with respect to the electrolyte.
Electrochemical cell which electrodes are Zinc and Copper submerged on Zinc sulfate and Copper sulfate respectively is known as Daniells cell.
Half reactions for a Daniells cell are these:
In order to avoid positive charges accumulating on the anode's compartment, an inverted U—shaped tube filled with an electrolytic solution is placed on the cell, thus allowing flow of electrons, producing D.C.
A voltameter is capable of measuring the change of electrical potential between the anode and the cathode.
Electrochemical cell voltage is also referred to as electromotive force or emf.
A cell diagram can be used to trace the path of the electrons in the electrochemical cell. For example, here is a cell diagram of a Daniells cell:
First, the reduced form of the metal to be oxidized at the anode (Zn) is written .
Standard electrode potential
Standard electrode potential is the value of the standard emf of a cell in which molecular hydrogen under standard pressure (105 Pa) is oxidized to solvated protons at the left-hand electrode.
The cell potential depends on the difference between each half cell potential.
Standard cell potential is calculated by the difference between the standard reduction potentials of each electrode.
It is impossible to measure directly half cell standard reduction potential, to avoid this problem a standard reduction potential is assignated to a reference acting as an electrode equivalent to .
The standard hydrogen electrode or (SHE) consists on an inverted glass tube similar to a laboratory test tube, where a light and fine platinum wire is connected to a thin platinum blade.
SHE operates exactly as the same way as conventional electrodes on Daniells cell's work;
For example on Copper standard reduction potential:
At standard temperature pressure conditions cell's emf (measured by a multimeter) is 0.34 V, conventionally SHE has a zero value, thus replacing on previous equation gives:
Electrochemical cell's emf value is used to predict whether redox reaction is a spontaneous process or not.
Changes over stoichiometric coefficients on balanced cell equation will not change value because standard electrode potential are intensive properties.
Spontaneity of Redox systems
On electrochemical cells, chemical energy transforms into electrical energy and is expressed mathematically as the product between cell's emf by electrical charge in Coulombs.
Electrochemical cell's total charge is determined by multiplying the number of moles by Faraday's constant (F).
Faraday's constant is the electrical charge in 1 mole of electrons, it has been measured experimentally and is equivalent to 96 485.3 coulombs.
Cell's emf measured is the maximum voltage produced, this value is used to calculate the maximum electrical energy which is obtained from a chemical reaction, this energy is referred to as electrical work and is expressed on the following equation,
,thus free energy is the amount of mechanical (or other) work that can be extracted from a system, replacing this value on previous equation with gives the relation between spontaneity and electrochemical cells.
The relation between Gibbs free energy and maximum electrical work may predict (at standard temperature and pressure conditions) whether cell's redox system is a spontaneous process or not.
A spontaneous electrochemical reaction can be used to generate an electrical current, in electrochemical cells.
Conversely, non-spontaneous electrochemical reactions can be driven forward by the application of a current at sufficient voltage.
The relation between equilibrium constant and spontaneity based on Gibbs free energy terms on electrochemical cells is expressed as follows:
Solving both equations express cell's mathematical relation between standard potential, and equilibrium constant.
Previous equation can use Briggsian logarithm as shown below:
Cell emf dependency on changes in concentration
Nernst Equation
Calculating cell's potential is not always plausible at standard temperature and pressure conditions.
On mid 1800s Willard Gibbs formulated an equation for spontaneous process at any conditions,
,Willard stated Q's dependency over reactants and products activity and designated it as their respective chemical activity.
Walther based on Willard Gibbs work during the mid 19th century, formulated a new equation where replaced 's value with cell's respective maximum electrical work, on Gibbs equation.
Finally he replaced 's value with electrochemical cell potential, thus formulating a new equation which now bears his name.
Assuming standard conditions () and R = the equation above can be expressed on Base—10 logarithm as shown below:
Concentration cells
A concentration cell is an electrochemical cell whose electrodes are from the same material differing in ionic concentrations on both half-cells.
For example an electrochemical cell, where two copper electrodes are submerged on blue vitriol's solution, whose concentrations are 0.05 M and 2.0 M , while connected through wire and saline bridge.
Le Chatelier's principle indicates reaction is favourable to reduction as concentration of ions increases.
The following cell diagram describes the cell mentioned above:
Where both half cell reactions for oxidation and reduction are:
Where cell's emf is calculated through Nernst equation as follows:
's value of this kind of cell is zero, as electrodes and ions are the same in both half-cells. After replacing values from case mentioned is possible to calculate cell's potential:
However, this value is only approximate, because the potential difference is given from the ratio of activities of the ions, not the ratio of concentrations.
Concentration cell's are often a significant biologist's matter of investigation hence they are present on biological cells where membrane potential is responsible of nerve synapses and cardiac beat.
Battery
A battery is an electrochemical cell or a group of them, where if combined together, may produce direct current at a constant voltage.
Dry cell
Dry cells don't have a fluid electrolyte instead they use a moist electrolyte paste.
Leclanché's simplified half reactions are shown below:
The voltage obtained from the zinc-carbon battery is 1.5 V approximately.
Mercury battery
Mercury battery has many applications on medicine and electronics.
Mercury battery half reactions are shown below:
There are no changes on the electrolyte's composition when cell works.
Lead-acid battery
The Lead-acid battery used on automobiles, consists on a series of six identical cells in line assembled, each cell has a lead anode and a cathode made from lead dioxide packed in a metal plaque.
Lead-acid battery half cell reactions are shown below:
At standard conditions, each cell may produce a direct current of 2 V, hence overall voltage produced is 12 V.
Solid state Lithium battery
Most of the batteries work using an aqueous electrolyte or a moist electrolyte paste instead, however a solid state battery operates using a solid electrolyte.
Flow battery/ Redox flow battery
Most batteries have all of the electrolyte and electrodes within a single housing.
These types of batteries are typically used for large-scale energy storage (kWh - multi MWh).
Fuel cells
Fossil fuels are used on power plants to supply electrical needs of a certain area, however the conversion of them into electricity is a low efficient process, in fact the most efficient electrical power plant it may convert into electricity about 40% of the original chemical energy when burned or processed.
To enhance electrical production, scientists developed fuel cells where combustion reactions are stimulated by electrochemical methods, thus requiring continuous replenishment of the reactants consumed.
The most popular is the oxygen-hydrogen fuel cell, where two inert–electrodes (porous electrodes of Nickel and Nickel oxide) are placed in an electrolytic solution such as hot caustic potash, in both compartments (anode and cathode) gaseous hydrogen and oxygen are bubbled into solution.
Oxygen-hydrogen fuel cell reactions are shown bellow:
The overall reaction is some-like to hydrogen combustion, differing on oxidation and reduction took place in anode and cathode separately, similar to the electrode used in the cell for measuring standard reduction potential having a double function acting as electrical conductors providing a surface required to decomposition of the molecules into atoms before electron transferring, thus named electrocatalysts.
Corrosion
Corrosion is the term applied to metal rust caused by an electrochemical process.
Iron corrosion
For iron rust to occur the metal has to be in contact with oxygen and water, although chemical reactions for this process are relatively complex and not all of them are completely understood, it is believed the causes are the following:
Electron transferring (Reduction-Oxidation) One area on the surface of the metal acts as the anode, which is where the oxidation (corrosion) occurs. H ions oxides, following this equation:Iron(III) oxide hydrated is known as rust.
Corrosion of coinage metals
Coinage metals, such as copper and silver, can also slowly corrode.
Prevention of Corrosion
Attempts to save a metal from becoming anodic are of two general types.
While it is almost impossible to prevent anode/cathode formation, if a non-conducting material covers the metal contact with the electrolyte is not possible and corrosion will not occur.
Coating
Metals are coated on its surface with paint or some other non-conducting coating.
Other prevention is called passivation where a metal is coated with another metal such as tin can.
Sacrificial anodes
A method commonly used to protect a structural metal is to attach a metal which is more anodic than the metal to be protected.
Electrolysis
Spontaneous redox reactions produces electricity, thus passage of electrons through a wire in the electric circuit.
Electrolysis of molten sodium chloride
When molten, sodium chloride can be electrolysed to yield metallic sodium and gaseous chlorine.
Reactions that take place at Down's cell are the following:
This process can yield industrial amounts of metallic sodium and gaseous chlorine, and is widely used on mineral dressing and metallurgy industries.
Standard emf for this process is approximately -4 V indicating a non-spontaneous process.
Electrolysis of water
Water at standard temperature and pressure conditions doesn't decompose into hydrogen and oxygen spontaneously as the Gibbs free energy for the process at standard conditions is about 474.4 kJ
However, special laboratory glassware has been designed for this purpose- the Hofmann voltameter. The following half reactions describe the process mentioned above:
Although strong acids may be used in the apparatus, the reaction will not net consume the acid.
Electrolysis of aqueous solutions
Electrolysis in an aqueous is a similar process as mentioned in electrolysis of water.
Electrolysis of a solution of Sodium chloride
The presence of water in a solution of sodium chloride must be examined in respect to its reduction and oxidation in both electrodes.
The following half reactions describes the process mentioned:
Reaction 1 is discarded as it has the most negative value on standard reduction potential thus making it less thermodynamically favorable in the process.
When comparing the reduction potentials in reactions 2 &
Finally, reaction 3 is favorable because it describes the proliferation of OH ions less favorable an option.
The overall reaction for the process according to the analysis would be the following:
As the overall reaction indicates, the concentration of chloride ions is reduced in comparison to OH- ions (whose concentration increases).
First law
Faraday concluded after several experiments on electrical current in non-spontaneous process, the mass of the products yielded on the electrodes was proportional to the value of current supplied to the cell, the length of time the current existed, and the molar mass of the substance analyzed.
In other words, the amount of a substance deposited on each electrode of an electrolytic cell is directly proportional to the quantity of electricity passed through the cell.
Below a simplified equation of Faraday's first law:
Where,
m is the mass of the substance produced at the electrode (in grams), Q is the total electric charge that passed through the solution (in coulombs), n is the valence number of the substance as an ion in solution (electrons per ion), M is the molar mass of the substance (in grams per mole).An important aspect of the second law of electrolysis is electroplating which together with the first law of electrolysis, has a significant number of applications in the industry, as when used to protect metals to avoid corrosion.
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