Electrochemistry is a branch of
chemistrythat studies chemical reactionswhich take place in a solutionat the interface of an electron conductor (a metalor a semiconductor) and an ionic conductor (the electrolyte), and which involve electron transfer between the electrode and the electrolyte or species in solution.
If a chemical reaction is driven by an external applied
voltage, as in electrolysis, or if a voltage is created by a chemical reaction as in a battery, it is an "electrochemical" reaction. Chemical reactions where electrons are transferred between molecules are called oxidation/reduction ( redox) reactions. In general, electrochemistry deals with situations where oxidationand reduction reactions are separated in space or time, connected by an external electric circuit to understand each process.
16th to 18th century developments
The 16th century marked the beginning of electrical understanding. During that century the English scientist
William Gilbertspent 17 years experimenting with magnetismand, 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 Guerickecreated the first electric generator, which produced static electricity by applying friction in the machine. The generator was made of a large sulfurball cast inside a glass globe, mounted on a shaft. The ball was rotated by means of a crank and a static electric sparkwas 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—18th century the French
chemist Charles-Augustin de Coulombdeveloped the law of electrostaticattraction in 1781 as an outgrowth of his attempt to investigate the law of electrical repulsions as stated by Joseph Priestleyin England.In the late 18th century the Italian physicianand anatomist Luigi Galvanimarked 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.
In 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 lightningor by the electric eeland torpedo ray as well as the "artificial" form produced by friction(i.e., static electricity).
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.
In 1800, William Nicholson and
Johann Wilhelm Rittersucceeded in decomposing water into hydrogenand oxygenby electrolysis. Soon thereafter Ritter discovered the process of electroplating. He also observed that the amount of metal deposited and the amount of oxygen produced during an electrolytic process 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 Wollastonmade 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 sodiumand potassiumfrom their compounds and of the alkaline earth metalsfrom 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 electromagnetismto others. André-Marie Ampèrequickly repeated Ørsted's experiment, and formulated them mathematically.
In 1821, Estonian-German
physicist Thomas Johann Seebeckdemonstrated the electrical potential in the juncture points of two dissimilar metals when there is a heatdifference between the joints.
In 1827 the German scientist
Georg Ohmexpressed 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.
Michael Faraday's experiments led him to state his two laws of electrochemistry. In 1836 John Daniell invented a primary cell in which hydrogenwas eliminated in the generation of the electricity. Daniell had solved the problem of polarization. In his laboratory he had learned that alloying the amalgamated zincof Sturgeon with mercury would produce a better voltage.
William Grove produced the first
fuel cellin 1839. In 1846, Wilhelm Weberdeveloped the electrodynamometer. In 1866, Georges Leclanchépatented a new cell which eventually became the forerunner to the world's first widely used battery, the zinc carbon cell. Svante August Arrheniuspublished his thesis in 1884 on "Recherches sur la conductibilité galvanique des électrolytes" (Investigations on the galvanic conductivity of electrolytes). 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.
Paul Héroultand Charles M. Hall developed a successful method to obtain aluminiumby using the principles described by Michael Faraday.
In 1894 Friedrich Ostwald concluded important studies of the
electrical conductivityand electrolytic dissociation of organic acids.
Walther Hermann Nernst developed the theory of the
electromotive forceof 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.
Fritz Habershowed that definite reduction products can result from electrolytic processes if the potential at the cathodeis kept constant. In 1898 he explained the reduction of nitrobenzenein stages at the cathode and this became the model for other similar reduction processes.
The 20th century and recent developments
The Electrochemical Society(ECS) was founded.
Robert Andrews Millikanbegan a series of experiments to determine the electric charge carried by a single electron.
Johannes Nicolaus Brønstedand Thomas Martin Lowrypublished essentially the same theory about how acids and bases behave, using an electrochemical basis. Arne Tiseliusdeveloped the first sophisticated electrophoreticapparatus in 1937 and some years later he was awarded the 1948 Nobel Prizefor his work in protein electrophoresis.
A year later, in 1949, the
International Society of Electrochemistry(ISE) was founded.
By the 1960s–1970s
quantum electrochemistrywas developed by Revaz Dogonadzeand his pupils.
Electrochemical processes involve
redoxreactions where an electronis transferred to or from a moleculeor ionchanging its oxidation state. This reaction can occur through the application of an external voltageor through the release of chemical energy.
Oxidation and reduction
The atoms, ions, or molecules involved in an electrochemical reaction are characterized by the number of
electrons each has compared to its number of protons called its " oxidation state" and is denoted by a + or a -. Thus the superoxide ion, O2-, has an "oxidation state" of -1. An atom or ion that gives up an electron to another atom or ion has its oxidation state increase, and the recipient of the negatively charged electron has its oxidation state decrease. Oxidation and reduction always occur in a paired fashion such that one species is oxidized when another is reduced. This paired electron transfer is called a redoxreaction.
For example when atomic
sodiumreacts with atomic chlorine, sodium donates one electron and attains an oxidation state of +1. Chlorine accepts the electron and its oxidation state is reduced to −1. The sign of the oxidation state (positive/negative) 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 from an atom or molecule is called
oxidation, and the gain of electrons is reduction. This can be easily remembered through the use of mnemonicdevices. Two of the most popular are "OIL RIG" (Oxidation Is Loss, Reduction Is Gain) and "LEO" the lion says "GER" (Lose Electrons: Oxidization, Gain Electrons: Reduction). For cases where electrons are shared (covalent bonds) between atoms with large differences in electronegativity, the electron is assigned to the atom with the largest electronegativity in determining the oxidation state.
The atom or molecule which loses electrons is 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. Oxygen is a common oxidizing agent, but not the only one. Despite the name, an oxidation reaction does not necessarily need to involve oxygen. In fact, a
firecan be fed by an oxidant other than oxygen; fluorinefires are often unquenchable, as fluorine is an even stronger oxidant (it has a higher electronegativity) than oxygen.
For reactions involving oxygen, the gain of oxygen implies the oxidation of the atom or molecule to which the oxygen is added (and the oxygen is reduced). For example, in the oxidation of
octaneby oxygento form carbon dioxideand water, both the carbon in the octane and the oxygen begin with an oxidation state of 0. In forming CO2 the carbon loses four electrons to become C4+ and the oxygens each gain two electrons to be O2-. In organic compounds, such as butaneor ethanol, the loss of hydrogen implies oxidation of the molecule from which it is lost (and the hydrogen is reduced). This follows because the hydrogen donates its electron in covalent bonds with non-metals but it takes the electron along when it is lost. Conversely, loss of oxygen or gain of hydrogen implies reduction.
Balancing redox reactions
Electrochemical reactions in water are better understood by balancing redox reactions using the
Ion-Electron Methodwhere H+ , OH- ion, H2O and electrons (to compensate the oxidation changes) are added to cell's half reactions for oxidation and reduction.
In acid medium H+ ions and water are added to
half reactions to balance the overall reaction.For example, when manganesereacts with 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::
In basic medium OH- ions and water are added to half reactions to balance the overall reaction. For example on reaction between
Potassium permanganateand 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::
The same procedure as used on acid medium is applied, for example on balancing using electron ion method to complete combustion of
propane.:::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::
An electrochemical cell is a device that produces an electric current from energy released by a spontaneous redox reaction. This kind of cell includes the
Galvanic cellor Voltaic cell, named after Luigi Galvaniand Alessandro Volta, both scientists who conducted several experiments on chemical reactions and electric current during the late 18th century.
Electrochemical cells have two conductive electrodes (the anode and the cathode). The
anodeis defined as the electrode where oxidation occurs and the cathodeis the electrode where the reduction takes place. Electrodes can be made from any sufficiently conductive materials, such as metals, semiconductors, graphite, and even conductive polymers. In between these electrodes is the electrolyte, which contains ions that can freely move.
The Galvanic cell uses two different metal electrodes, each in an electrolyte where the positively charged ions are the oxidized form of the electrode metal. One electrode will undergo oxidation (the anode) and the other will undergo reduction (the cathode). The metal of the anode will oxidize, going from an oxidation state of 0 (in the solid form) to a positive oxidation state and become an ion. At the cathode, the metal ion in solution will accept one or more electrons from the cathode and the ion's oxidation state is reduced to 0. This forms a solid metal that electrodeposits on the cathode. The two electrodes must be electrically connected to each other, allowing for a flow of electrons that leave the metal of the anode and flow through this connection to the ions at the surface of the cathode. This flow of electrons is an electrical current that can be used to do work, such as turn a motor or power a light.
A Galvanic cell whose
electrodes are zincand coppersubmerged in zinc sulfateand copper sulfate, respectively, is known as a Daniell cell.
Half reactions for a Daniell cell are these:::
In this example, the anode is zinc metal which oxidizes (loses electrons) to form zinc ions in solution, and copper ions accept electrons from the copper metal electrode and the ions deposit at the copper cathode as an electrodeposit. This cell forms a simple battery as it will spontaneously generate a flow of electrical current from the anode to the cathode through the external connection. This reaction can be driven in reverse by applying a voltage, resulting in the deposition of zinc metal at the anode and formation of copper ions at the cathode. To provide a complete electric circuit, there must also be an ionic conduction path between the anode and cathode electrolytes in addition to the electron conduction path. The simplest ionic conduction path is to provide a liquid junction. To avoid mixing between the two electrolytes, the liquid junction can be provided through a porous plug that allows ion flow while reducing electrolyte mixing. To further minimize mixing of the electrolytes, a
salt bridgecan be used which consists of an electrolyte saturated gel in an inverted U-tube. As the negatively charged electrons flow in one direction around this circuit, the positively charged metal ions flow in the opposite direction in the electrolyte.
Electrochemical cell voltage is also referred to as
electromotive forceor 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 Daniell cell::First, the reduced form of the metal to be oxidized at the anode (Zn) is written. This is separated from its oxidized form by a vertical line, which represents the limit between the phases (oxidation changes). 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. The electrolyte concentration is given as it is an important variable in determining the cell potential.
tandard electrode potential
To allow prediction of the cell potential, tabulations of
standard electrode potentialare available. Such tabulations are referenced to the standard hydrogen electrode (SHE). The standard hydrogen electrodeundergoes the reaction:which is shown as reduction but, in fact, the SHE can act as either the anode or the cathode, depending on the relative oxidation/reduction potential of the other electrode/electrolyte combination. The term standard in SHE requires a supply of hydrogen gas bubbled through the electrolyte at a pressure of 1 atm and an acidic electrolyte with H+ activity equal to 1 (usually assumed to be [H+] = 1 mol/liter).
The SHE electrode can be connected to any other electrode by a salt bridge to form a cell. If the second electrode is also at standard conditions, then the measured cell potential is called the standard electrode potential for the electrode. The standard electrode potential for the SHE is zero, by definition. The polarity of the standard electrode potential provides information about the relative reduction potential of the electrode compared to the SHE. If the electrode has a positive potential with respect to the SHE, then that means it is a strongly reducing electrode which forces the SHE to be the anode (an example is Cu in aqueous CuSO4 with a standard electrode potential of 0.337 V). Conversely, if the measured potential is negative, the electrode is more oxidizing than the SHE (such as Zn in ZnSO4 where the standard electrode potential is -0.763 V).
Standard electrode potentials are usually tabulated as reduction potentials. However, the reactions are reversible and the role of a particular electrode in a cell depends on the relative oxidation/reduction potential of both electrodes. The oxidation potential for a particular electrode is just the negative of the reduction potential. A standard cell potential can be determined by looking up the standard electrode potentials for both electrodes (sometimes called half cell potentials). The one that is smaller will be the anode and will undergo oxidation. The cell potential is then calculated as the sum of the reduction potential for the cathode and the oxidation potential for the anode.
For example, the standard electrode potential for a copper electrode is:
:::At standard temperature, pressure and concentration conditions, the cell's emf (measured by a
multimeter) is 0.34 V. by definition, the electrode potential for the SHE is zero. Thus, the Cu is the cathode and the SHE is the anode giving:Or,:
Changes in the
stoichiometric coefficients of a balanced cell equation will not change value because the standard electrode potential is an intensive property.
pontaneity of Redox reaction
During operation of electrochemical cells,
chemical energyis transformed into electrical energyand is expressed mathematically as the product of the cell's emf and the electrical chargetransferred through the external circuit.:where is the cell potential measured in volts (V) and is the cell current integrated over time and measured in coulumbs (C). can also be determined by multiplying the total number of electrons transferred (measured in moles) times Faraday's constant, F = 96,485 C/mole.
The emf of the cell at zero current is the maximum possible emf. It is used to calculate the maximum possible electrical energy that could be obtained from a
chemical reaction. This energy is referred to as electrical workand is expressed by the following equation:
:where work is defined as positive into the system.
Since the free energy is the maximum amount of work that can be extracted from a system, one can write: :
A positive cell potential gives a negative change in Gibbs free energy. This is consistent with the cell production of an electric current flowing from the cathode to the anode through the external circuit. If the current is driven in the opposite direction by imposing an external potential, then work is done on the cell to drive electrolysis.
spontaneouselectrochemical reaction (change in Gibbs free energy less than zero) can be used to generate an electric current, in electrochemical cells. This is the basis of all batteries and fuel cells. For example, gaseous oxygen (O2) andhydrogen (H2) can be combined in a fuel cell to form water andenergy, typically a combination of heat and electrical energy.
Conversely, non-spontaneous electrochemical reactions can be driven forward by the application of a current at sufficient
voltage. The electrolysisof water into gaseous oxygen and hydrogen is a typical example.
The relation between the
equilibrium constant, "K", and the Gibbs free energy for an electrochemical cell is expressed as follows:
Rearranging to express the relation between standard potential and equilibrium constant yields
:Previous equation can use
Briggsian logarithmas shown below::
Cell emf dependency on changes in concentration
The standard potential of an electrochemical cell requires standard conditions for all of the reactants. When reactant concentrations differ from standard conditions, the cell potential will deviate from the standard potential. In the 20th century German
chemist Walther Hermann Nernstproposed a mathematical model to determine the effect of reactant concentration on electrochemical cell potential.
In the late 19th century
Josiah Willard Gibbshad formulated a theory to predict whether a chemical reaction is spontaneous based on the free energy
"ΔG" = change in
Gibbs free energy, "T" = absolute temperature, "R" = gas constant, ln = natural logarithm, "Q" = reaction quotient.
Gibbs' key contribution was to formalize the understanding of the effect of reactant concentration on spontaneity.
Based on Gibbs' work, Nernst extended the theory to include the contribution from electric potential on charged species. As shown in the previous section, the change in Gibbs free energy for an electrochemical cell can be related to the cell potential. Thus, Gibbs' theory becomes
Finally, Nernst divided through by the amount of charge transferred to arrive at 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::
A concentration cell is an electrochemical cell where the two electrodes are the same material, the electrolytes on the two half-cells involve the same ions, but the electrolyte concentration differs between the two half-cells.
For example an electrochemical cell, where two copper electrodes are submerged in two
copper(II) sulfatesolutions, whose concentrations are 0.05 M and 2.0 M, connected through a salt bridge. This type of cell will generate a potential that can be predicted by the Nernst equation. Both electrodes undergo the same chemistry (although the reaction proceeds in reverse at the cathode)
Le Chatelier's principleindicates that the reaction is more favourable to reduction as the concentration of ions increases. Reduction will take place in the cell's compartment where concentration is higher and oxidation will occur on the more dilute side.
The following cell diagram describes the cell mentioned above::Where the half cell reactions for oxidation and reduction are::::
Where the 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 the case mentioned, it is possible to calculate cell's potential::
However, this value is only approximate, as reaction quotient is defined in terms of ion activities which can be approximated with the concentrations as calculated here.
The Nernst equation plays an important role in understanding electrical effects in cells and organelles. Such effects include nerve
synapsesand cardiac beat as well as the resting potential of a somatic cell.
A battery is an electrochemical cell (sometimes several in series) used for chemical energy storage. Batteries are optimized to produce a constant electric current for as long as possible. Although the cells discussed previously are useful for theoretical purposes and some laboratory experiments, the large internal resistance of the salt bridge make them inappropriate battery technologies. Various alternative battery technologies have been commercialized as discussed next.
Dry cells do not have a
fluidelectrolyte. Instead, they use a moist electrolyte paste. Leclanché's cell is a good example of this, where the anode is a zinccontainer surrounded by a thin layer of manganese dioxideand a moist electrolyte paste of ammonium chlorideand zinc chloridemixed with starch. The cell's cathode is represented by a carbon bar inserted on the cell's electrolyte, usually placed in the middle.
Leclanché's simplified half reactions are shown below:::::
The voltage obtained from the
zinc-carbon batteryis around 1.5 V.
The mercury battery has many applications in
medicineand electronics. The battery consists of a steel—made container in 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 oxideand mercury(II) oxide.
Mercury battery half reactions are shown below:::::There are no changes in the electrolyte's composition when the cell works. Such batteries provide 1.35 V of
The lead-acid battery used in
automobiles, consists of a series of six identical cells assembled in series. Each cell has a leadanode and a cathode made from lead dioxidepacked in a metalplaque. Cathode and anode are submerged in a solution of sulfuric acidacting as the electrolyte.
Lead-acid battery half cell reactions are shown below:::
At standard conditions, each cell may produce a potential of 2 V, hence overall voltage produced is 12 V. Differing from mercury and zinc-carbon batteries, lead-acid batteries are rechargeable. If an external voltage is supplied to the battery it will produce an
electrolysisof the products in the overall reaction (discharge), thus recovering initial components which made the battery work.
Lithium rechargeable battery
Instead of an
aqueouselectrolyte or a moist electrolyte paste, a solid state battery operates using a solid electrolyte. Lithiumpolymer batteries are an example of this; a graphite bar acts as the anode, a bar of lithium cobaltate acts as the cathode, and a polymer, swollen with a lithium salt, allows the passage of ionsand serves as the electrolyte. In this cell, the carbon in the anode can reversibly form a lithium-carbon alloy. Upon discharging, lithium ions spontaneously leave the lithium cobaltate cathode and travel through the polymer and into the carbon anode forming the alloy. This flow of positive lithium ions is the electrical current that the battery provides. By charging the cell, the lithium dealloys and travels back into the cathode. The advantage of this kind of battery is that Lithium possess the highest negative value of standard reduction potential. It is also a light metaland therefore less mass is required to generate 1 mole ofelectrons. Lithium ion battery technologies are widely used in portable electronic devices because they have high energy storage density and are rechargeable. These technologies show promise for future automotive applications, with new materials such as iron phosphates and lithium vanadates.
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 (kWh - multi MWh). Of the several different types that have been developed, some are of current commercial interest, including the
vanadium redox batteryand zinc bromine battery.
Fossil fuelsare used in power plantsto supply electrical needs, however their conversion into electricity is an inefficient process. The most efficient electrical power plant may only convert about 40% of the original chemical energyinto electricity when burned or processed.
To enhance electrical production, scientists have developed fuel cells where
combustionis replaced by electrochemical methods, similar to a battery but requiring continuous replenishment of the reactantsconsumed.
The most popular is the oxygen-hydrogen fuel cell, where two inert electrodes (
porouselectrodes of nickeland nickel oxide) are placed in an electrolytic solution such as hot caustic potash, in both compartments (anode and cathode) gaseous hydrogenand oxygenare bubbled into solution.
Oxygen-hydrogen fuel cell reactions are shown bellow::::
The overall reaction is identical to hydrogen
combustion. Oxidation and reduction take place in the anode and cathodeseparately. This is similar to the electrode used in the cell for measuring standard reduction potential which has a double function acting as electrical conductorsproviding a surface required to decomposition of the moleculesinto atomsbefore electron transferring, thus named electrocatalysts. Platinum, nickel, and rhodiumare good electrocatalysts.
Corrosion is the term applied to
metal rustcaused by an electrochemical process. Most people are likely familiar with the corrosion of iron, in the form of reddish rust. Other examples include the black tarnish on silver, and red or green corrosion that may appear on copperand its alloys, such as brass. The cost of replacing metals lost to corrosion is in the multi-billions of dollars per year.
For iron rust to occur the metal has to be in contact with
oxygenand 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. At the anode, the metal gives up electrons.
Electronsare transferred from ironreducing oxygen in the atmosphereinto water on the cathode, which is placed in another region of the metal.
##Global reaction for the process:
emffor iron rusting:
###:Iron corrosion takes place on acid medium; H+
ionscome from reaction between carbon dioxidein the atmosphere and water, forming carbonic acid. Fe2+ ions oxides, following this equation:: Iron(III) oxide hydratedis known as rust. The concentration of water associated with iron oxide varies, thus chemical representation is presented as . The electric circuitworks as passage of electrons and ions occurs, thus if an electrolyte is present it will facilitate oxidation, this explains why rusting is quicker on salt water.
Corrosion of common metals
Coinage metals, such as copper and silver, slowly corrode through use.A patinaof green-blue copper carbonateforms on the surface of copperwith exposure to the water and carbon dioxide in the air. Silvercoins or cutlerythat are exposed to high sulfur foods such as eggs or the low levels of sulfur species in the air develop a layer of black Silver sulfide. Goldand platinumare extremely difficult to oxidize under normal circumstances, and require exposure to a powerful chemical oxidizing agent such as aqua regia.
Some common metals oxidize extremely rapidly in air.
Titaniumand aluminium oxidize instantaneouly in contact with the oxygen in the air. These metals form an extremely thin layer of oxidized metal on the surface. This thin layer of oxide protects the underlying layers of the metal from the air preventing the entire metal from oxidizing. These metals are used in applications where corrosion resistance is important. Iron, in contrast, has an oxide that forms in air and water, called rust, that does not stop the further oxidation of the iron. Thus iron left exposed to air and water will continue to rust until all of the iron is oxidized.
Prevention of corrosion
Attempts to save a metal from becoming anodic are of two general types. Anodic regions dissolve and destroy the structural integrity of the metal.
While it is almost impossible to prevent anode/
cathodeformation, if a non-conducting material covers the metal, contact with the electrolyteis not possible and corrosion will not occur.
coated on its surface with paintor some other non-conducting coating. This prevents the electrolytefrom reaching the metal surface if the coating is complete. Scratches exposing the metal will corrode with the region under the paint, adjacent to the scratch, to be anodic.
Other prevention is called "
passivation" where a metal is coated with another metal such as a tin can. Tin is a metal that rapidly corrodes to form a mono-molecular oxidecoating that prevents further corrosion of the tin. The tin prevents the electrolyte from reaching the base metal, usually steel( iron). However, if the tin coating is scratched the iron becomes anodic and can corrodes rapidly.
A method commonly used to protect a structural metal is to attach a metal which is more anodic than the metal to be protected. This forces the structural metal to be
cathodic, thus spared corrosion. It is called "sacrificial" because the anode dissolves and has to be replaced periodically. Zincbars are attached at various locations on steel shiphulls to render the ship hull cathodic. The zinc bars are replaced periodically. Other metals, such as magnesium, would work very well but zinc is the least expensive useful metal.
To protect pipelines, an ingot of buried or exposed magnesium (or zinc) is buried beside the pipeline and is connected electrically to the pipe above ground. The pipeline is forced to be a cathode and is protected from being oxidized and rusting. The magnesium anode is sacrificed. At intervals new
ingots are buried to replace those lost.
The spontaneous redox reactions of a conventional battery produce electricity through the different chemical potentials of the cathode and anode in the electrolyte. However, electrolysis requires an external source of
electrical energyto induce a chemical reaction, and this process takes place in a compartment called electrolytic cell.
Electrolysis of molten sodium chloride
When molten, the salt
sodium chloridecan be electrolyzed to yield metallic sodiumand gaseous chlorine. Industrially this process takes place in a special cell named Down's cell. The cell is connected to an electrical power supply, allowing electronsto migrate from the power supply to the electrolytic cell.
Reactions that take place at Down's cell are the following::::
This process can yield large amounts of metallic sodium and gaseous chlorine, and is widely used on
mineral dressingand metallurgyindustries.
emffor this process is approximately -4 Vindicating a (very) non-spontaneous process. In order for this reaction to occur the power supply should provide at least a potential of 4 V. However, larger voltages must be used for this reaction to occur at a high rate.
Electrolysis of water
Water can be converted to its component elemental gasses, H2 and O2 through the application of an external voltage.
Waterdoesn't decompose into hydrogenand oxygenspontaneously as the Gibbs free energyfor the process at standard conditions is about 474.4 kJ. The decomposition of water into hydrogen and oxygen can be performed in an electrolytic cell. In it, a pair of inert electrodesusually made of platinumimmersed in water act as anode and cathode in the electrolytic process. The electrolysis starts with the application of an external voltage between the electrodes. This process will not occur except at extremely high voltages without an electrolyte such as sodium chlorideor sulfuric acid(most used 0.1 M).
Bubbles from the gases will be seen near both electrodes. 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. While this reaction will work at any conductive electrode at a sufficiently large potential, platinum catalyzes both hydrogen and oxygen formation, allowing for relatively mild voltages (~2V depending on the
Electrolysis of aqueous solutions
Electrolysis in an aqueous is a similar process as mentioned in electrolysis of water. However, it is considered to be a complex process because the contents in solution have to be analyzed in half reactions, whether reduced or oxidized.
Electrolysis of a solution of sodium chloride
The presence of water in a solution of
sodium chloridemust be examined in respect to its reduction and oxidation in both electrodes. Usually, water is electrolysed as mentioned in electrolysis of water yielding "gaseous oxygenin the anode" and gaseous hydrogenin the cathode. On the other hand, sodium chloride in water dissociates in Na+ and Cl- ions, anionwill be attracted to the cathode, thus reducing the sodiumion. The cationwill then be attracted to the anode oxidizing chlorideion.
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 & 4, the reduction of chloride ion is favored. Thus, if the Cl- ion is favored for reduction, then the water reaction is favored for
oxidationproducing gaseous oxygen, however experiments shown gaseous chlorine is produced and not oxygen.
Although the initial analysis is correct, there is another effect that can happen, known as the overvoltage effect. Additional voltage is sometimes required, beyond the voltage predicted by the . This may be due to kinetic rather than thermodynamic considerations. In fact, it has been proven that the
activation energyfor the chloride ion is very low, hence favorable in kinetic terms. In other words, although the voltage applied is thermodynamically sufficient to drive electrolysis, the rate is so slow that to make the process proceed in a reasonable time frame, the voltageof the external source has to be increased (hence, overvoltage).
The overall reaction for the process according to the analysis would be the following::::
As the overall reaction indicates, the
concentrationof chloride ions is reduced in comparison to OH- ions (whose concentration increases). The reaction also shows the production of gaseous hydrogen, chlorineand aqueous sodium hydroxide.
Quantitative electrolysis & Faraday's Laws
Quantitative aspects of electrolysis were originally developed by
Michael Faradayin 1834. Faraday is also credited to have coined the terms " electrolyte", electrolysis, among many others while he studied quantitative analysis of electrochemical reactions. Also he was an advocate of the law of conservation of energy.
Faraday concluded after several experiments on
electrical currentin non-spontaneous process, the massof 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 electricitypassed 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).
Faraday devised the laws of chemical electrodeposition of metals from solutions in 1857. He formulated the second law of electrolysis stating "the amounts of bodies which are equivalent to each other in their ordinary chemical action have equal quantities of electricity naturally associated with them." In other terms, the quantities of different elements deposited by a given amount of electricity are in the
ratioof their chemical equivalent weights.
An important aspect of the second law of electrolysis is
electroplatingwhich 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.
There are various extremely important electrochemical processes in both nature and industry, like the coating of objects with metals or metal oxides through electrodeposition and the detection of alcohol in drunken drivers through the redox reaction of ethanol. The generation of chemical energy through
photosynthesisis inherently an electrochemical process, as is production of metals like aluminum and titanium from their ores. Certain diabetes blood sugar meters measure the amount of glucose in the blood through its redox potential.
The nervous impulses in neurons are based on electric current generated by the movement of sodium and potassium ions into and out of cells, and certain animals like eels can generate a powerful voltage from certain cells that can disable much larger animals.
Activity series of metals
Contact tension- a historical forerunner to the theory of electrochemistry.
* Important publications in electrochemistry
Table of standard electrode potentials
* cite web
title = Electrochemistry
work = General Chemistry II by Dr. Michael Blaber
url = http://wine1.sb.fsu.edu/chm1046/notes/Electro/Electrol/Electrol.htm
accessmonthday = January 30
accessyear = 2006
* cite web
title = Electrochemistry
work = Corrosion
url = http://colossus.chem.umass.edu/chandler/ch112/corrosion.htm
accessmonthday = January 28
accessyear = 2006
* cite web
title = Michael Faraday
work = Biography
url = http://scienceworld.wolfram.com/biography/Faraday.html
accessmonthday = January 30
accessyear = 2006
* cite web
title = The Faraday law of electrochemistry
work = Faraday laws of electrochemistry
url = http://library.thinkquest.org/19662/low/eng/electrolysis.html
accessmonthday = January 30
accessyear = 2006
* cite book
first = Raymond
last = Chang
authorlink = Raymond Chang
year = 2002
title = Chemistry
chapter = Electrochemistry
edition = 7th Edition
publisher = Mc Graw Hill
isbn = 0-07-365601-1
url = http://www.mhhe.com/chang7
oclc = 45393784
* cite book
first = Theodore
last = L. Brown
authorlink = Theodore L. Brown
coauthors = H. Eugene LeMay, Jr., Bruce E. Bursten, Julia R. Burdge
year = 2003
title = Chemistry
chapter = Electrochemistry
edition = 9th Edition
publisher = Pearson Education
location = US
isbn = 0-13-066997-0
url = http://www.pearsoneducation.net/brown
oclc = 47953855
* cite book
first = John
last = William Hill
authorlink = John William Hill
coauthors = Ralph H. Petrucci, Terry McCreary, Scott S. Perry
year = 2004
month = March
title = General Chemistry: An Integrated Approach
chapter = Electrochemistry
edition = 7th Edition
pages = 1200
publisher = Pearson Education
isbn = 0-13-140283-8
* cite book
first = John
last = McMurry
authorlink = John McMurry
coauthors = Robert C. Fay
year = 2004
month = March
title = Chemistry
chapter = Electrochemistry
edition = 3rd Edition
publisher = Prentice Hall
isbn = 0-13-056765-5
url = http://cwx.prenhall.com/bookbind/pubbooks/mcmurry2/
first = Keith
last = Laidler
authorlink = Keith James Laidler
coauthors = John H. Meiser, Bryan C. Sanctuary
year = 2002
month = May
title = Physical Chemistry
chapter = Electrochemistry
editor = Keith Laidler
edition = 4th Edition
publisher = Houghton Mifflin Company College Division
Boston, Massachusetts, U.S.
id = ISBN 061815292X
url = http://college.hmco.com/CollegeCatalog/CatalogController?cmd=Portal&subcmd=display&ProductID=59
first = A.J.
last = Bard
coauthors = Inzelt, G.; Scholz F. (eds)
year = 2008
title = Electrochemical Dictionary
url = http://www.springer.com/chemistry/physical/book/978-3-540-74597-6
pages = 723 pp, more than 2770 entries
publisher = Springer
id = isbn = 978-3-540-74597-6
* [http://www.electrochemistry.net Electrochemistry net (commercial link)]
* [http://www.electrochem.org The Electrochemical Society]
* [http://electrochem.cwru.edu/estir/ Electrochemical Science and Technology Information Resource (ESTIR) ]
* [http://www.ise-online.org International Society of Electrochemistry (ISE)]
* [http://electrochem.cwru.edu/ed/encycl/ Electrochemistry Encyclopedia at Case Western Reserve University]
* [http://electrochem.cwru.edu/ed/dict.htm Electrochemistry Dictionary at Case Western Reserve University] (size ~ 388KB)
* [http://www.funsci.com/fun3_en/electro/electro.htm Experiments in Electrochemistry at Fun Science]
* [http://bip.cnrs-mrs.fr/bip06/software.html SOAS: a free (open-source) software to analyze electrochemical data]
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