Encyclopedia
Electrochemistry is a branch of
chemistry that studies the reactions which take place at the interface of an electronic conductor and an ionic conductor .
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. In general, electrochemistry deals with situations where an
oxidation and a
reduction reaction is separated in space. The direct charge transfer from one molecule to another is not the topic of electrochemistry.
History
16th to 18th century developments
The
16th century marked the beginning of the electrical understanding. During the 1550s the English scientist
William Gilbert spent 17 years experimenting with
magnetism and, to a lesser extent, electricity. For his work on magnets, Gilbert became known as the
"Father of Magnetism." He discovered various methods for producing and strengthening magnets.
In 1663 the
German physicist Otto von Guericke created the first electric generator, which produced static electricity by applying friction in the machine. The generator was made of a large
sulfur ball cast inside a glass globe, mounted on a shaft. The ball was rotated by means of a crank and a static electric
spark was produced when a pad was rubbed against the ball as it rotated. The globe could be removed and used as source for experiments with electricity.
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. Du Fay announced that electricity consisted of two fluids:
"vitreous" , or positive, electricity; and
"resinous," or negative, electricity. This was the
two-fluid theory of electricity, which was to be opposed by
Benjamin Franklin's one-fluid theory later in the century.
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" 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. He believed that this new force was a form of electricity in addition to the
"natural" form produced by
lightning or by the
electric eel and
torpedo ray as well as the
"artificial" form produced by friction .
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. Galvani refuted this by obtaining muscular action with two pieces of the same material.
19th century
In 1800, the English chemists William Nicholson and Johann Ritter succeeded in decomposing water into
hydrogen and
oxygen by
electrolysis. Soon thereafter Johann Ritter discovered the process of
electroplating. He also observed the amount of metal deposited and the amount of oxygen produced during an electrolytic process that depended on the distance between the electrodes. By 1801 Ritter observed thermoelectric currents and anticipated the discovery of thermoelectricity by Thomas Johann Seebeck.
By the 1810s
William Hyde Wollaston made improvements to the galvanic pile.
Sir
Humphry Davy's work with electrolysis led to the conclusion that the production of electricity in simple electrolytic cells resulted from chemical action and that chemical combination occurred between substances of opposite charge. This work led directly to the isolation of
sodium and
potassium from their compounds and of the alkaline earth metals from theirs in 1808.
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.
André-Marie Ampère quickly repeated Ørsted's experiment, and formulated them mathematically.
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" 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. In 1836
John Daniell invented a primary cell in which
hydrogen was eliminated in the generation of the electricity. Daniell had solved the problem of polarization. In his laboratory he had learned that
alloying the amalgamated
zinc of Sturgeon with mercury would produce a better voltage.
William Grove produced the first
fuel cell in 1839. In 1846,
Wilhelm Weber developed the
electrodynamometer. In 1866, Georges Leclanché patented a new cell which eventually became the forerunner to the world's first widely used battery, the
zinc carbon cell.
Svante August Arrhenius published his thesis in 1884 on
Recherches sur la conductibilité galvanique des électrolytes . From his results the author concluded that electrolytes, when dissolved in water, become to varying degrees split or dissociated into electrically opposite positive and negative ions.
In 1886
Paul Héroult and
Charles M. Hall developed a successful method to obtain
aluminum by using the principles described by Michael Faraday.
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 1889, he showed how the characteristics of the current produced could be used to calculate the free energy change in the chemical reaction producing the current. He constructed an equation, known as Nernst Equation, which related the voltage of a cell to its properties.
In 1898
Fritz Haber showed that definite reduction products can result from electrolytic processes if the potential at the
cathode is kept constant. In 1898 he explained the reduction of
nitrobenzene in stages at the cathode and this became the model for other similar reduction processes.
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.
Arne Tiselius developed the first sophisticated
electrophoretic apparatus in 1937
and some years later he was awarded to the 1948
Nobel Prize for his work in protein
electrophoresis.
A year later, in 1949, the International Society of Electrochemistry was founded.
By the
1960s–
1970s quantum electrochemistry was developed by
Revaz Dogonadze and his pupils.
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.
In a redox reaction, an atom's oxidation state changes as a result of an electron transfer.
Oxidation and Reduction
The
elements involved in an electrochemical
reaction are characterized by the number of
electrons each has. The
oxidation state of an ion is the number of electrons it has accepted or donated compared to its neutral state . If an
atom or ion donates an
electron in a reaction its oxidation state is increased, if an element accepts an electron its oxidation state is decreased.
For example when
sodium reacts with chlorine, sodium donates one electron and gains an oxidation state of +1. Chlorine accepts the electron and gains an oxidation state of -1. The sign of the oxidation state actually corresponds to the value of each ion's electronic charge. The attraction of the differently charged sodium and chlorine ions is the reason they then form an
ionic bond.
The loss of electrons of a substance is called
oxidation, and the gain of electrons is
reduction. This can be easily remembered through the use of
mnemonic devices. Two of the most popular are
"OIL RIG" and
"LEO" the lion says
"GER" .
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 oxidizing agent is always being reduced in a reaction; the reducing agent is always being oxidized.
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. These are very common; as one substance loses electrons the other substance accepts them.
Oxidation requires an oxidant. Oxygen is an oxidant, but not the only one. Despite the name, an oxidation reaction does not necessarily need to involve oxygen. In fact, even
fire can be fed by an oxidant other than oxygen:
fluorine fires are often unquenchable, as fluorine is an even stronger oxidant than oxygen.
Balancing redox reactions
Electrochemical reactions in water are better understood by balancing redox reactions using the Ion-Electron Method where
H+ , OH
- ion,
H2O and electrons 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.
For example on
Manganese reacts to Sodium bismuthate.
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. For example on reaction between
Potassium permanganate and
Sodium sulfite.
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. This kind of cell is also known as
Galvanic cell or Voltaic cell, named after
Luigi Galvani and
Alessandro Volta, both scientists conducted several experiments on chemical reactions and electric current during the late
18th century.
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. Each metal in the Galvanic cell undergoes a different half-reaction. This causes the metals to have different dissolving rates, leading to an unequal number of electrons in the two metals. This results in a different electrode potential between the electrolyte and each metal. If an electrical connection, such as a
wire or direct contact, is formed between the two, an electric current flows between the metals.
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. electric current.
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 is written . This is separated from its oxidised form by a vertical line, which represents the limit between the phases . The double vertical lines represent the saline bridge on the cell. Finally, the oxidized form of the metal to be reduced at the cathode, is written, separated from its reduced form by the vertical line.
Standard electrode potential
Standard electrode potential is the value of the standard emf of a cell in which molecular hydrogen under standard pressure is oxidized to solvated protons at the left-hand electrode.
The cell potential depends on the difference between each half cell potential. Conventionally the potential associated with each electrode is chosen as the
reduction takes place on the chosen electrode, hence standard electrode potential are tabulated on reduction potentials, thus tables are built on standard reduction potentials noted as .
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 . Cell's half reaction used for this procedure is
hydrogen which in standard temperature and pressure conditions acts as a zero volt electrode.
The
standard hydrogen electrode or 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. This setup is placed in a solution of
Hydrochloric acid, plenty of H
+ ions, gaseous
hydrogen enter through the tube and react over the platinum blade thus allowing reduction and oxidation processes to occur.
SHE operates exactly as the same way as conventional electrodes on Daniells cell's work; in order to measure the standard reduction potential, SHE replaces one of the electrodes in the electrochemical cell acting as
cathode or
anode, thus electric current generated on the cell represents the standard reduction potential for the element which is measured.
For example on Copper standard reduction potential:
At standard temperature pressure conditions cell's emf 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. A positive sign for overall cell's standard potential is considered to be spontaneous reaction, a negative sign would predict a spontaneous reaction on the opposite direction.
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 .
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 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 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. This is the basis of all batteries and
fuel cells. For example, gaseous oxygen and
hydrogen can be combined in a fuel cell to form water and
energy .
Conversely, non-spontaneous electrochemical reactions can be driven forward by the application of a current at sufficient
voltage. The
electrolysis of water into gaseous oxygen and hydrogen is a typical example.
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. However in
1900s German
chemist Walther Hermann Nernst proposed a mathematical model to determine electrochemical cell potential where standard conditions cannot be reached.
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. Reduction will take place in cell's compartment where concentration is higher and oxidation will occur on the diluted side.
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. Electrochemical principles which made batteries work are the same as on electrochemical cells, however a battery doesn't need auxiliary components such as saline bridge on Daniell cells.
Dry cell
Dry cells don't have a fluid electrolyte instead they use a moist electrolyte paste.
Leclanché's cell is a good example of this, where cell's
anode is a
zinc container surrounded by a thin layer of
manganese dioxide and a moist electrolyte paste of
ammonium chloride and
zinc chloride mixed with starch to have a pale and flabby consistency and avoiding flees. Cell's cathode is represented by a carbon bar inserted on cell's electrolyte, usually placed in the middle.
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. The battery consists on a
steel—made container with the shape of a cylinder acting as the cathode, where an amalgamated anode of mercury and zinc is surrounded by a stronger alkaline electrolyte and a paste of
Zinc oxide and
Mercury oxide .
Mercury battery half reactions are shown below:
There are no changes on the electrolyte's composition when cell works. Mercurium battery provides 1.35 V of
direct current.
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. Cathode and anode are submerged in a solution of
sulfuric acid acting as the electrolyte.
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. Lead-acid batteries, differing from Mercury and Zinc-carbon batteries, are
rechargeable. If an external voltage is supplied to the battery it will produce an
electrolysis of the products in the overall reaction , thus recovering initial components which made the battery work.
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. Solid state
lithium batteries are an example of this, where a solid Lithium bar acts as the
anode, a bar of Lithium sulfide or Vanadium oxide acts as the
cathode and a
polymer, allowing the passage of ions and not
electrons, serves as the electrolyte. The advantage of this kind of battery from others is that Lithium possess the highest negative value of standard reduction potential. It is also a light metal and therefore less mass is required to generate 1 mole of electrons. This battery is rechargeable and it can provide a
direct current of about 3
V. Although solid state batteries are frowned upon nowadays, it is likely they will someday become a reliable source of
electricity.
Flow battery/ Redox flow battery
Most batteries have all of the electrolyte and electrodes within a single housing. A flow battery is unusual in that the majority of the electrolyte, including dissolved reactive species, is stored in separate tanks. The electrolytes are pumped through a reactor, which houses the electrodes, when the battery is charged or discharged.
These types of batteries are typically used for large-scale energy storage . Of the several different types that have been developed, some are of current commercial interest, including the
vanadium redox battery and zinc bromine battery.
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 are placed in an electrolytic solution such as hot
caustic potash, in both compartments 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.
Platinum,
nickel,
rhodium are good electrocatalysts.
Corrosion
Corrosion is the term applied to
metal rust caused by an electrochemical process. Most people are likely familiar with the corrosion of
iron, in the form or reddish rust. Other examples include the black tarnish on
