Electrochemistry

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'''Electrochemistry''' is a branch of [[chemistry]] that studies the reactions which take place at the interface of an electronic [[Electrical conductor|conductor]] (the [[electrode]] composed of a [[metal]] or a [[semiconductor]], including [[graphite]]) and an ionic conductor (the [[electrolyte]]).

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If a [[chemical reaction]] is caused by an external [[voltage]], or if a voltage is caused by a chemical reaction, as in a [[battery (electricity)|battery]], it is an ''electrochemical'' reaction. In general, electrochemistry deals with situations where an [[oxidation]] and a [[redox|reduction]] reaction are separated in space. The direct [[charge transfer]] from one molecule to another is not the topic of electrochemistry.

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==History==

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===16th to 18th century developments===

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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.

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In [[1663]] the [[Germany|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 electricity|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.

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By the mid—[[1700s]] the [[France|French]] [[chemist]] [[C.F. du Fay|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"'' (from the [[Latin language|Latin]] for ''"glass"''), or positive, electricity; and ''"resinous,"'' or negative, electricity. This was the ''two-fluid theory'' of electricity, which was to be opposed by [[Benjamin Franklin|Benjamin Franklin's]] ''one-fluid theory'' later in the century.[[Image:Galvani-frog-legs.PNG|thumb|left|200px|Late [[1780s]] diagram of Galvani's experiment on frog legs.]]

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[[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.

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In the late [[1700s]] the [[Italy|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.

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On his essay Galvani concluded that animal tissue contained a here-to-fore neglected innate, vital force, which he termed ''"animal electricity,"'' which activated [[nerve]]s and [[muscle]]s spanned by [[metal]] [[probe]]s. 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 [[Electric ray|torpedo ray]] as well as the ''"artificial"'' form produced by [[friction]] (i.e., static electricity).

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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.

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===19th century===

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[[Image:Humphrydavy.jpg|thumb|left|150px|Sir Humphry Davy's portrait in [[1800s]].]]

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In [[1800]], the English chemists [[William Nicholson (chemist)]] 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 [[thermoelectricity|thermoelectric currents]] and anticipated the discovery of thermoelectricity by [[Thomas Johann Seebeck]].

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By the [[1810s]] [[William Hyde Wollaston]] made improvements to the [[galvanic pile]].

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Sir [[Humphry Davy]]'s work with electrolysis led to the conclusion that the production of electricity in simple [[electrolytic cell]]s 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]].

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[[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.

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[[Image:ChemicalHistoryofaCandle.PNG|thumb|right|140px|Professor Michael Faraday's portrait on his book [[The Chemical History of a Candle]].]]

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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.

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In [[1827]] the German scientist [[Georg Ohm]] expressed his [[Ohm's law|law]] in this famous book ''"Die galvanische Kette, mathematisch bearbeitet"'' (The Galvanic Circuit Investigated Mathematically) in which he gave his complete theory of electricity.

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In [[1832]] [[Michael Faraday]]'s experiments on Electrochemistry led him to state his two laws of electrochemistry. In [[1836]] [[John Frederic Daniell|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 [[alloy]]ing the [[amalgam|amalgamated]] [[zinc]] of Sturgeon with [[Mercury (element)|mercury]] would produce a better voltage.

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[[Image:Arrhenius2.jpg|thumb|left|140px|Swedish chemist [[Svante Arrhenius]] portrait circa [[1880s]].]]

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[[William Robert Grove|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 battery|zinc carbon cell]].

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[[Svante August Arrhenius]] published 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 [[electrolyte]]s, when dissolved in water, become to varying degrees split or dissociated into electrically opposite positive and negative ions.

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In [[1886]] [[Paul Héroult]] and [[Charles Martin Hall|Charles M. Hall]] developed a successful method to obtain [[aluminum]] by using the principles described by Michael Faraday.

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In [[1894]] [[Wilhelm Ostwald|Friedrich Ostwald]] concluded important studies of the [[electrical conductivity]] and electrolytic dissociation of [[organic acid]]s.

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[[Image:Walther Nernst 2.jpg|thumb|right|140px|German scientist [[Walther Nernst]] portrait in [[1910s]].]]

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[[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.

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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.

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===The 20th century and recent developments ===

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In [[1902]], [[The Electrochemical Society]] (ECS) was founded.

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In [[1909]], [[Robert Andrews Millikan]] began a series of experiments to determine the electric charge carried by a single [[electron]].

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In [[1923]], [[Johannes Nicolaus Brønsted]] and [[Thomas Martin Lowry]] published essentially the same theory about how acids and bases behave, using an electrochemical basis.

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[[Arne Tiselius]] developed the first sophisticated [[electrophoretic]] apparatus in [[1937]]

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and some years later he was awarded to the [[1948]] [[Nobel Prize]] for his work in protein [[electrophoresis]].

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A year later, in [[1949]], the [[International Society of Electrochemistry]] (ISE) was founded.

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By the [[1960s]]–[[1970s]] [[quantum electrochemistry]] was developed by [[Revaz Dogonadze]] and his pupils.

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==Electrolysis==

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Spontaneous redox reactions produces electricity, thus passage of electrons through a wire in the [[electric circuit]]. Electrolysis requires an external source of [[electrical energy]] to induce a chemical reaction, this process takes place in a compartment called [[electrolytic cell]]. Principles involved on electrolysis are the same as featured on electrochemical cells.

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===Electrolysis of molten sodium chloride===

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When molten, [[sodium chloride]] can be electrolysed to yield metallic [[sodium]] and gaseous [[chlorine]]. Industrially this process takes place in a special cell named Down's cell. The cell is connected to a battery, allowing [[electrons]] [[to migrate|migration]] from the battery to the electrolytic cell.

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Reactions that take place at Down's cell are the following:

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:<math>\mbox{Anode (oxidation): }2Cl^{-} \rightarrow Cl_{2}(g) + 2e^{-}\,</math>

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:<math>\mbox{Cathode (reduction): }2Na^{+}(l) + 2e^{-} \rightarrow 2Na(l)\,</math>

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:<math>\mbox{Overall reaction: }2Na^{+} + 2Cl^{-}(l) \rightarrow 2Na(l) + Cl_{2}(g)\,</math>

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This process can yield industrial amounts of metallic sodium and gaseous chlorine, and is widely used on [[mineral dressing]] and [[metallurgy]] [[industry|industries]].

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Standard [[emf]] for this process is approximately -4 [[V]] indicating a non-spontaneous process. In order this reaction to occur the battery should provide at least a potential of 4V. However, on mineral refining industry, higher voltages are used, due to low efficiency on the process.

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===Electrolysis of water===

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[[Image:Hoffman voltameter.jpg|thumb|190px|Diagram of a Hofmann voltameter, showing electrolysis of water.]]

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[[Water]] at standard temperature and pressure conditions doesn't decompose into [[hydrogen]] and [[oxygen]] [[Spontaneous process|spontaneously]] as the [[Gibbs free energy]] for the process at standard conditions is about 474.4 kJ

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However, special [[laboratory glassware]] has been designed for this purpose- the [[Hofmann voltameter]]. In it, a pair of inert [[electrodes]] usually made of [[platinum]] act as anode and cathode in the electrolytic process. After the water (if pure) has been placed in the [[apparatus]], nothing happens, hence there are not enough [[ions]] to let the passage of electrons occur. To start the electrolysis an electrolyte should be placed in, usually [[sodium chloride]] or [[sulfuric acid]] (most used 0.1 [[Molar concentration|M]]).

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Bubbles from the gases will be seen near both electrodes. The following half reactions describe the process mentioned above:

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:<math>\mbox{Anode (oxidation): }2H_{2}O(l) \rightarrow O_{2}(g) + 4H^{+}(aq) + 4e^{-}\,</math>

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:<math>\mbox{Cathode (reduction): }2H_{2}O(g) + 2e^{-} \rightarrow H_{2}(g) + 2OH^{-}(aq)\,</math>

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:<math>\mbox{Overall reaction: }2H_{2}O(l) \rightarrow 2H_{2}(g) + O_{2}(g)\,</math>

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Although strong acids may be used in the apparatus, the reaction will not net consume the acid.

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===Electrolysis of aqueous solutions===

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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 [[chemical reaction|half reactions]], whether reduced or oxidized.

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====Electrolysis of a solution of Sodium chloride====

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The presence of water in a solution of [[sodium chloride]] must be examined in respect to its reduction and oxidation in both electrodes. Usually, water is electrolysed as mentioned in electrolysis of water yielding ''gaseous [[oxygen]] in the anode'' and gaseous [[hydrogen]] in the cathode. On the other hand, sodium chloride in water [[Dissociation (chemistry)|dissociates]] in Na+ and Cl- ions, [[anion]] will be attracted to the cathode, thus reducing the [[sodium]] ion. The [[cation]] will then be attracted to the anode oxidizing [[chloride]] ion.

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The following half reactions describes the process mentioned:

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:<math>\mbox{1. Cathode: }Na^{+}(aq)+ 1e^{-} \rightarrow Na(s) \qquad E^{o}_{red}=-2.71 V\,</math>

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:<math>\mbox{2. Anode: }2Cl^{-}(aq) \rightarrow Cl_{2}(g) + 2e^{-} \qquad E^{o}_{red}= +1.36 V\,</math>

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:<math>\mbox{3. Cathode: }2H_{2}O(l) + 2e^{+} \rightarrow H_{2}(g) + 2OH^{-}(aq)\qquad E^{o}_{red}=-0.83 V\,</math>

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:<math>\mbox{4. Anode: } 2H_{2}O(l) \rightarrow O_{2}(g) + 4H^{+}(aq) + 4e^{-}\qquad E^{o}_{red}=+1.23V\,</math>

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Reaction 1 is discarded as it has the most [[Negative and non-negative numbers|negative]] value on standard reduction potential thus making it less thermodynamically favorable in the process.

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When comparing the reduction potentials in reactions 2 & 4, the reduction of chloride ion is favored. Thus, if the Cl- ion is favored for [[redox|reduction]], then the water reaction is favored for [[oxidation]] producing gaseous oxygen, however experiments shown gaseous chlorine is produced and not oxygen.

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Although the initial analysis is correct, there is another effect that can happen, known as the [[Overvoltage|overvoltage effect]]. Additional voltage is sometimes required, beyond the voltage predicted by the <math>E^{o}_{cell}\,</math>. This may be due to [[chemical kinetics|kinetic]] rather than [[Thermochemistry|thermodynamic]] considerations. In fact, it has been proven that the [[activation energy]] for the chloride ion is very low, hence favorable in [[chemical kinetics|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 [[voltage]] of the external source has to be increased (hence, overvoltage).

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Finally, reaction 3 is favorable because it describes the proliferation of [[Hydroxide|OH-]] ions thus letting a probable reduction of [[Proton|H+]] ions less favorable an option.

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The overall reaction for the process according to the analysis would be the following:

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:<math>\mbox{Anode (Oxidation): } 2Cl^{-}(aq)\rightarrow Cl_{2}(g) + 2e^{-}\,</math>

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:<math>\mbox{Cathode (Reduction): } 2H_{2}O(l) + 2e{-}\rightarrow H_{2}(g) + 2OH^{-}(aq)\,</math>

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:<math>\mbox{Overall reaction: } 2H_{2}O + 2Cl^{-}(aq) \rightarrow H_{2}(g) + Cl_{2}(g) + 2OH^{-}(aq)\,</math>

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As the overall reaction indicates, the [[concentration]] of chloride ions is reduced in comparison to OH- ions (whose concentration increases). The reaction also shows the production of gaseous [[hydrogen]], [[chlorine]] and aqueous [[sodium hydroxide]].

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===Quantitative electrolysis & Faraday Laws===

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Quantitative aspects of electrolysis were originally developed by [[Michael Faraday]] in [[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]].

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[[Image:Nickel-electroplating.PNG|thumb|200px|right|To [[refining|refine]] metals a process named [[electroplating]] is used (diagram shows [[Nickel]] refining); the process has its bases on the first and the second law of electrolysis stated by Faraday in the [[19th century]].]]

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====First law====

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Faraday concluded after several experiments on [[electrical current]] in [[spontaneous process|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.

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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.

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Below a simplified equation of Faraday's first law:

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:<math>m \ = \ { 1 \over F \ } \cdot { Q M \over n } </math>

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Where,

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:''m'' is the mass of the substance produced at the electrode (in [[grams]]),

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:''Q'' is the total electric charge that passed through the solution (in [[coulomb]]s),

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:''n'' is the valence number of the substance as an ion in solution (electrons per ion),

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:''F'' is the Faraday's constant = 96,485 [[coulomb]]s per mole,

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:''M'' is the molar mass of the substance (in grams per [[mole (unit)|mole]]).

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Faraday's constant is equal to the amount of charge carried by one mole of electrons.

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====Second law====

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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 [[ratio]] of their chemical [[equivalent weight]]s.

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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 [[metal]]s to avoid corrosion.

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==Principles==

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===Redox reactions===

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Electrochemical process are redox reactions where [[energy]] is produced by a [[Spontaneous process|spontaneous reaction]] which produces electricity, otherwise [[electrical current]] stimulates a chemical reaction.

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In a redox reaction, an atom's oxidation state changes as a result of an [[electron transfer]].

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===Oxidation and Reduction===

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The [[chemical element|element]]s involved in an electrochemical [[chemical reaction|reaction]] are characterized by the number of [[electron]]s each has. The ''oxidation state'' of an [[ion]] is the number of electrons it has accepted or donated compared to its neutral state (which is defined as having an oxidation state of 0). 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.

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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 (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]].

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The loss of electrons of a substance is called [[oxidation]], and the gain of electrons is [[redox|reduction]]. This can be easily remembered through the use of [[mnemonic]] devices. 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).

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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.

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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.

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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.

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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 (it has a higher [[electronegativity]]) than oxygen.

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===Balancing redox reactions===

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Electrochemical reactions in water are better understood by balancing redox reactions using the [[Ion-Electron Method]] where [[Proton|H+]] , [[Hydroxide|OH-]] ion, [[Water (molecule)|H2O]] and electrons (to compensate the oxidation changes) are added to cell's [[half reaction]]s for oxidation and reduction.

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====Acid medium====

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In acid medium [[Proton|H]] atoms and water are added to [[half reaction]]s to balance the overall reaction.

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For example, when [[Manganese]] reacts with [[Sodium bismuthate]].

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:<math>\mbox{Reaction unbalanced: }\mbox{Mn}^{2+}(aq) + \mbox{NaBiO}_3(s)\rightarrow\mbox{Bi}^{3+}(aq) + \mbox{MnO}_4^{-}(aq)\,</math>

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:<math>\mbox{Oxidation: }\mbox{4H}_2\mbox{O}(l)+\mbox{Mn}^{2+}(aq)\rightarrow\mbox{MnO}_4^{-}(aq) + \mbox{8H}^{+}(aq)+\mbox{5e}^{-}\,</math>

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:<math>\mbox{Reduction: }\mbox{2e}^{-}+ \mbox{6H}^{+}(aq) + \mbox{BiO}_3^{-}(s)\rightarrow\mbox{Bi}^{3+}(aq) + \mbox{3H}_2\mbox{O}(l)\,</math>

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Finally the reaction is balanced by [[multiplication|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.

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:<math>\mbox{8H}_2\mbox{O}(l)+\mbox{2Mn}^{2+}(aq)\rightarrow\mbox{2MnO}_4^{-}(aq) + \mbox{16H}^{+}(aq)+\mbox{10e}^{-}\,</math>

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:<math>\mbox{10e}^{-}+ \mbox{30H}^{+}(aq) + \mbox{5BiO}_3^{-}(s)\rightarrow\mbox{5Bi}^{3+}(aq) + \mbox{15H}_2\mbox{O}(l)\,</math>

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Reaction balanced:

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:<math>\mbox{14H}^{+}(aq) + \mbox{2Mn}^{2+}(aq)+ \mbox{5NaBiO}_3(s)\rightarrow\mbox{7H}_2\mbox{O}(l) + \mbox{2MnO}_4^{-}(aq)+\mbox{5Bi}^{3+}(aq)+\mbox{5Na}^{+}(aq)\,</math>

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====Basic medium====

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In basic medium [[Hydroxide|OH-]] ions and [[Water (molecule)|water]] are added to half reactions to balance the overall reaction. For example on reaction between [[Potassium permanganate]] and [[Sodium sulfite]].

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:<math>\mbox{Reaction unbalanced: }\mbox{KMnO}_{4}+\mbox{Na}_{2}\mbox{SO}_3+\mbox{H}_2\mbox{O}\rightarrow\mbox{MnO}_{2}+\mbox{Na}_{2}\mbox{SO}_{4}+\mbox{KOH}\,</math>

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:<math>\mbox{Reduction: }\mbox{3e}^{-}+\mbox{2H}_{2}\mbox{O}+\mbox{MnO}_{4}^{-}\rightarrow\mbox{MnO}_{2}+\mbox{4OH}^{-}\,</math>

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:<math>\mbox{Oxidation: }\mbox{2OH}^{-}+\mbox{SO}^{2-}_{3}\rightarrow\mbox{SO}^{2-}_{4}+\mbox{H}_{2}\mbox{O}+\mbox{2e}^{-}\,</math>

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The same procedure as followed on acid medium by multiplying electrons to opposite half reactions solve the equation thus balancing the overall reaction.

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:<math>\mbox{6e}^{-}+\mbox{4H}_{2}\mbox{O}+\mbox{2MnO}_{4}^{-}\rightarrow\mbox{2MnO}_{2}+\mbox{8OH}^{-}\,</math>

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:<math>\mbox{6OH}^{-}+\mbox{3SO}^{2-}_{3}\rightarrow\mbox{3SO}^{2-}_{4}+\mbox{3H}_{2}\mbox{O}+\mbox{6e}^{-}\,</math>

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Equation balanced:

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:<math>\mbox{2KMnO}_{4}+\mbox{3Na}_{2}\mbox{SO}_3+\mbox{H}_2\mbox{O}\rightarrow\mbox{2MnO}_{2}+\mbox{3Na}_{2}\mbox{SO}_{4}+\mbox{2KOH}\,</math>

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====Neutral medium====

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The same procedure as used on acid medium is applied, for example on balancing using electron ion method to [[Combustion|complete combustion]] of [[propane gas]].

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:<math>\mbox{Reaction unbalanced: }\mbox{C}_{3}\mbox{H}_{8}+\mbox{O}_{2}\rightarrow\mbox{CO}_{2}+\mbox{H}_{2}\mbox{O}\,</math>

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:<math>\mbox{Reduction: }\mbox{4H}^{+} + \mbox{O}_{2}\rightarrow\mbox{H}_{2}\mbox{O}+\mbox{H}_{2}\mbox{O}+ \mbox{4e}^{-}\,</math>

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:<math>\mbox{Oxidation: }\mbox{20e}^{-}+\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\mbox{20H}^{+}\,</math>

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As in acid and basic medium, electrons which were used to compensate oxidation changes are multiplied to opposite half reactions, thus solving the equation.

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:<math>\mbox{20H}^{+}+\mbox{5O}_{2}\rightarrow\mbox{5H}_{2}\mbox{O}+\mbox{5H}_{2}\mbox{O}+\mbox{20e}^{-}\,</math>

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:<math>\mbox{20e}^{-}+\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\mbox{20H}^{+}\,</math>

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Equation balanced:

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:<math>\mbox{C}_{3}\mbox{H}_{8}+\mbox{5O}_{2}\rightarrow\mbox{3CO}_{2}+\mbox{4H}_{2}\mbox{O}\,</math>

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==Electrochemical cells==

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An electrochemical cell is a device capable of producing electric current from energy released by a [[Spontaneous process|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 who conducted several experiments on chemical reactions and electric current during the late [[18th century]].

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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.

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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|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.

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An electrochemical cell which [[electrode]]s are [[Zinc]] and [[Copper]] submerged on [[Zinc sulfate]] and [[Copper sulfate]] respectively is known as [[Daniell cell|Daniells cell]].

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Half reactions for a Daniells cell are these:

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:<math>\mbox{Zinc electrode (anode) : }\mbox{Zn}(s)\rightarrow\mbox{Zn}^{2+}(aq)+\mbox{2e}^{-}\,</math>

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:<math>\mbox{Copper electrode (cathode) : }\mbox{Cu}^{2+}(aq)+\mbox{2e}^{-}\rightarrow\mbox{Cu}(s)\,</math>

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[[Image:BASi epsilon C3 cell stand.jpg|thumb|right||260px|A modern cell stand for electrochemical research. The electrodes attach to high-quality metallic wires, and the stand is attached to a potentiostat/galvanostat (not pictured). A [[shotglass]]-shaped container is [[Aerated water|aerated]] with a noble gas and sealed with the [[teflon]] block.]]

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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 [[Direct current|D.C.]] electric current.

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A [[galvanometer|voltameter]] is capable of measuring the change of [[Electric potential|electrical potential]] between the anode and the cathode.

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Electrochemical cell voltage is also referred to as [[electromotive force]] or [[emf]].

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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:

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:<math>\mbox{Zn}(s)|\mbox{Zn}^{2+}(1M)||\mbox{Cu}^{2+}(1M)|\mbox{Cu}(s)\,</math>

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First, the reduced form of the metal to be oxidized at the anode (Zn) is written . This is separated from its oxidised 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.

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==Standard electrode potential==

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[[Image:Reduction-potentials2.PNG|thumb|right|260px|The [[standard reduction potentials]] table is determined in a modified version of [[galvanic cell]] using an [[Hydrogen]] [[electrode]] as [[cathode]], because Hydrogen is taken as reference, standard reduction potential for that substance is zero (gray [[highlighter|highlight]]).]]

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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.

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The cell potential depends on the difference between each half cell potential. Conventionally the potential associated with each electrode is chosen as the [[redox|reduction]] takes place on the chosen electrode, hence standard electrode potential are [[tabulation|tabulated]] on reduction potentials, thus tables are built on [[standard reduction potential]]s noted as <math>\mbox{E}^{0}_{red}\,</math>.

−

Standard cell potential is calculated by the difference between the standard reduction potentials of each electrode.

−

:<math>\mbox{E}^{o}_{cell}=\mbox{E}^{o}_{red}(cathode)-\mbox{E}^{o}_{red}(anode)</math>

−

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 <math>\mbox{E}^{0}_{red}=0\,</math>. Cell's half reaction used for this procedure is [[hydrogen]] which in [[Standard conditions for temperature and pressure|standard temperature and pressure]] conditions (105 Pa, 298.15 K, 1 mol. L-1) acts as a zero volt electrode.

−

The [[standard hydrogen electrode]] or ([[Standard hydrogen electrode|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]]. 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.

−

[[Standard hydrogen electrode|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:

−

:<math>\mbox{Cell diagram}\,</math>

−

:<math>\mbox{Pt}(s)|\mbox{H}_{2}(1 atm)|\mbox{H}^{+}(1 M)||\mbox{Cu}^{2+}(1 M)|\mbox{Cu}(s)\,</math>

−

:<math>\mbox{E}^{o}_{cell}=\mbox{E}^{o}_{red}(cathode)-\mbox{E}^{o}_{red}(anode)</math>

−

At standard temperature pressure conditions cell's [[electromotive force|emf]] (measured by a [[multimeter]]) is 0.34 V, conventionally [[Standard hydrogen electrode|SHE]] has a zero value, thus replacing on previous equation gives:

−

:<math>\mbox{0.34V}_{cell}=\mbox{E}^{o}_{\mbox{Cu}^{2+}/\mbox{Cu}}-\mbox{E}^{o}_{\mbox{H}^{+}/\mbox{H}_{2}}</math>

−

:<math>\mbox{0.34V}_{cell}=\mbox{E}^{o}_{\mbox{Cu}^{2+}/\mbox{Cu}}-0</math>

−

Electrochemical cell's [[electromotive force|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 coefficient]]s on balanced cell equation will not change <math>\mbox{E}^{0}_{red}\,</math> value because standard electrode potential are [[Intensive and extensive properties|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 [[Coulomb]]s.

−

:<math>\mbox{Electrical energy}=(\mbox{volts})(\mbox{coulombs})\,</math>

−

:<math>\mbox{Electrical energy}=\mbox{joules}\,</math>

−

Electrochemical cell's total charge is determined by multiplying the number of moles by [[Faraday's constant]] (F).

−

:<math>\mbox{Total charge}=\mbox{n}\mbox{F}\,</math>

−

Faraday's constant is the electrical charge in 1 [[Mole (unit)|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,

−

:<math>\mbox{W}_{max}=\mbox{W}_{electrical}\,</math>

−

:<math>\mbox{W}_{max}=-\mbox{nFE}_{cell}\,</math>

−

,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 <math>\Delta G\,</math>gives the relation between spontaneity and electrochemical cells.

−

:<math>\Delta G=-\mbox{nFE}_{cell}\,</math>

−

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 (electricity)|current]], in [[electrochemical cell]]s. This is the basis of all [[battery (electricity)|batteries]] and [[fuel cell]]s. For example, gaseous oxygen (O2) and

−

hydrogen (H2) can be combined in a fuel cell to form water and

−

energy (a combination of heat and electrical energy, typically).

−

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:

−

:<math>\Delta G^{o}=\mbox{-RT ln K}\,</math>

−

:<math>\mbox{-nFE}^{o}_{cell}=\mbox{-RT ln K}\,</math>

−

Solving both equations express cell's mathematical relation between standard potential, and equilibrium constant.

−

:<math>\mbox{E}^{o}_{cell}={\mbox{RT} \over \mbox{nF}} \mbox{ln K}\,</math>

−

Previous equation can use [[Briggsian logarithm]] as shown below:

−

:<math>\mbox{E}^{o}_{cell}={0.0592 \mbox{V} \over \mbox{n}} \mbox{log K}\,</math>

−

==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.

−

In the mid [[1800s]] [[Willard Gibbs]] formulated an equation for spontaneous process at any conditions,

−

:<math>\Delta G=\Delta G^{o}+\mbox{RT ln Q}\,</math> ,

−

Where:

−

''ΔG'' = change in [[Gibbs free energy]], ''T'' = absolute [[temperature]], ''R'' = [[gas constant]], ln = [[natural logarithm]], ''Q'' = [[reaction quotient]].

−

Willard stated Q's dependency over reactants and products activity and designated it as their respective [[Activity (chemistry)|chemical activity]].

−

Walther based on Willard Gibbs work during the mid [[19th century]], formulated a new equation where replaced <math>\Delta G\,</math>'s value with cell's respective maximum electrical work, on Gibbs equation.

−

:<math>nF\Delta E = nF\Delta E^\circ - R T \ln Q \, \,</math>

−

Where:

−

''n'' = number of [[electrons]]/[[Mole (unit)|mole]] product, ''F'' = [[Faraday constant]] ([[coulomb]]s/[[Mole (unit)|mole]]), and ''ΔE'' = [[electrical potential of the reaction]].

−

Finally he replaced <math>-nF\Delta E\,</math>'s value with electrochemical cell potential, thus formulating a new equation which now bears his name.

−

:<math>\Delta E=\Delta E^{o}- {\mbox{RT} \over \mbox{nF}} \mbox{ln Q}\,</math>

−

Assuming standard conditions (<math>Temperature = 298 K , 25 C\,</math>) and [[Universal gas constant|R]] = <math>8.3145 {J \over K mol}</math> the equation above can be expressed on [[Common logarithm|Base—10 logarithm]] as shown below:

−

:<math>\Delta E=\Delta E^{o}- {\mbox{0.0592 V} \over \mbox{n}} \mbox{log Q}\,</math>

−

===Concentration cells===

−

[[Image:Cell-membrane-electrochemical.PNG|thumb|250px|Calculating [[membrane potential]] is good example where concentration cells are used in biology to understanding cell's [[metabolism]] such as [[Na-K pump|Na+(red) K+(blue) pump]].]]

−

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|blue vitriol's]] solution, whose concentrations are 0.05 [[Molar concentration|M]] and 2.0 [[Molar concentration|M]] , while connected through wire and saline bridge.

−

:<math>Cu^{2+}(aq)+2e^{-}\rightarrow \mbox{Cu}(s)</math>

−

[[Le Chatelier's principle]] indicates reaction is favourable to reduction as concentration of <math>Cu^{2+}\,</math> 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:

−

:<math>Cu(s)|Cu^{2+}(0.05 M)||Cu^{2+}(2.0 M)|Cu(s)\,</math>

−

Where both half cell reactions for oxidation and reduction are:

−

:<math>Oxidation: Cu(s)\rightarrow \mbox{Cu}^{2+} (0.05 M) + 2e^{-}\,</math>

−

:<math>Reduction: Cu^{2+} (2.0 M) +2e^{-} \rightarrow \mbox{Cu} (s)\,</math>

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:<math>Overall reaction: Cu^{2+} (2.0 M) \rightarrow \mbox{Cu}^{2+} (0.05 M)\,</math>

−

Where cell's emf is calculated through Nernst equation as follows:

−

:<math>E = E^{o}- {0.0257 V \over 2} ln {[Cu^{2+}]_{diluted}\over [Cu^{2+}]_{concentrated}}\,</math>

−

<math>E^{o}\,</math>'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:

−

:<math>E = 0- {0.0257 V \over 2} ln {0.05\over 2.0}\,</math>

−

:<math>E = 0.0474 V\,</math>

−

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 [[Synapses|nerve synapses]] and [[Cardiac cycle|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===

−

[[Image:Zincbattery.png|thumb|250px|Zinc carbon battery diagram.]]

−

Dry cells don't have a [[fluid]] electrolyte instead they use a moist electrolyte paste. [[Zinc-carbon battery|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.

−

[[Georges Leclanché|Leclanché's]] simplified half reactions are shown below:

−

:<math>Anode: Zn(s) \rightarrow Zn^{2+} (aq) + 2e^{-}\,</math>

−

:<math>Cathode: 2NH^{+}_{4}(aq)+ 2MnO_{2}(s) + 2e^{-}\rightarrow Mn_{2}O_{3}(s) + 2NH_{3} (aq) + H_{2}O (l)\,</math>

−

:<math>\mbox{Overall reaction:}\,</math>

−

:<math>Zn(s) + 2NH^{+}_{4}(aq)+ 2MnO_{2}(s) \rightarrow Zn^{2+}(aq) + Mn_{2}O_{3}(s) + 2NH_{3} (aq) + H_{2}O (l)\,</math>

−

The voltage obtained from the [[zinc-carbon battery]] is 1.5 [[Volt|V]] approximately.

−

===Mercury battery===

−

[[Image:Mercurybattery2.PNG|thumb|200px|Cutaway view of a Mercury battery diagram.]]

−

Mercury battery has many applications on [[medicine]] and [[electronics]]. The battery consists of a [[steel]]—made container with the shape of a cylinder acting as the cathode, where an [[amalgam|amalgamated]] anode of mercury and zinc is surrounded by a stronger alkaline electrolyte and a paste of [[Zinc oxide]] and [[Mercury(II) oxide]] .

−

Mercury battery half reactions are shown below:

−

:<math>Anode: Zn(Hg) + 2OH^{-} (aq) \rightarrow ZnO(s) + H_{2}O (l) + 2e^{-}\,</math>

−

:<math>Cathode: HgO(s) + H_{2}O(l) + 2e^{-}\rightarrow Hg(l) + 2OH^{-} (aq)\,</math>

−

:<math>\mbox{Overall reaction:}\,</math>

−

:<math>Zn(Hg) + HgO(s) \rightarrow ZnO(s) + Hg(l)\,</math>

−

There are no changes on the electrolyte's composition when cell works. Mercurium battery provides 1.35 V of [[direct current]].

−

===Lead-acid battery===

−

[[Image:Lead acid cell.jpg|thumb|100px|A sealed 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:

−

:<math>Anode: Pb(s) + SO^{2-}_{4}(aq) \rightarrow PbSO_{4}(s) + 2e^{-}\,</math>

−

:<math>Cathode: PbO_{2}(s) + 4H^{+}(aq) + SO^{2-}_{4}(aq) + 2e^{-} \rightarrow PbSO_{4}(s) + 2H_{2}O(l)\,</math>

−

<math>\mbox{Overall reaction:} Pb(s) + PbO_{2}(s) + 4H^{+}(aq)+2SO^{2-}_{4}(aq) \rightarrow 2PbSO_{4}(s) + 2H_{2}O(l)</math>

−

At standard conditions, each cell may produce a [[direct current]] of 2 [[Volts|V]], hence overall voltage produced is 12 V. Lead-acid batteries, differing from Mercury and Zinc-carbon batteries, are [[Rechargeable battery|rechargeable]]. If an external voltage is supplied to the battery it will produce an [[electrolysis]] of the products in the overall reaction (discharge), 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 [[faraday constant|mole of electrons]]. This battery is rechargeable and it can provide a [[direct current]] of about 3 [[Volts|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 (kWh - multi MWh). 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[[percentage|%]] of the original [[chemical energy]] when [[combustion|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 electrode|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:

−

:<math>Anode: 2H_{2}(g)+ 4OH^{-}(aq)\rightarrow 4H_{2}O(l)+4e^{-}\,</math>

−

:<math>Cathode: O_{2}(g)+ 2H_{2}O(l) + 4e^{-}\rightarrow 4OH^{-}(aq)\,</math>

−

:<math>\mbox{Overall reaction:} 2H_{2}(g) + O_{2}(g)\rightarrow 2H_{2}O(l)\,</math>

−

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 [[electrocatalyst]]s. [[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 of reddish rust. Other examples include the black tarnish on [[silver]], and red or green corrosion that may appear on [[copper]] and its alloys, such as [[brass]]. The cost of replacing metals lost to corrosion is in the multi-billions of [[American dollar|dollars]] per year.

−

===Iron corrosion===

−

[[Image:Iron-rusting-scheme.PNG|thumb|270px|right|Diagram showing a water [[drop|droplet]] over an iron surface. Electrochemical mechanisms involved develop iron rusting process.{{replacethisimage}}]]

−

For iron rust to occur the metal has to be in contact with [[oxygen]] and [[water]], although [[chemical reaction]]s 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.

−

###:<math>Fe(s)\rightarrow Fe^{2+}(aq) + 2e^{-}\,</math>

−

##[[Electrons]] are transferred from [[iron]] reducing oxygen in the [[atmosphere]] into [[water (molecule)|water]] on the cathode, which is placed in another region of the metal.

−

###:<math>O_{2}(g) + 4H^{+}(aq) + 4e^{-} \rightarrow 2H_{2}O(l)\,</math>

−

##Global reaction for the process:

−

##:<math>2Fe(s) + O_{2}(g) + 4H^{+}(aq) \rightarrow 2Fe^{2+}(aq) + 2H_{2}O(l)\,</math>

−

##Standard [[emf]] for iron rusting:

−

###:<math>E^{o}=E^{o}_{cathode}-E^{o}_{anode}\,</math>

−

###:<math>E^{o}=1.23V-(-0.44V)=1.67V\,</math>

−

Iron corrosion takes place on acid medium; [[Proton|H+]] [[ions]] come from reaction between [[carbon dioxide]] in the atmosphere and water, forming [[carbonic acid]]. Fe2+ ions oxides, following this equation:

−

:<math>4Fe^{2+}(aq) + O_{2}(g) + (4+2x)H_{2}O(l) \rightarrow 2Fe_{2}O_{3}.xH_{2}O + 8H^{+}(aq)</math>

−

[[Iron(III) oxide]] [[hydrated]] is known as rust. Water associated with iron oxide it varies, thus chemical representation is presented as <math>Fe_{2}O_{3}.xH_{2}O\,</math>.

−

The [[electric circuit]] works as passage of electrons and ions occurs, thus if an electrolyte is present it will facilitate [[oxidation]], this explains why rusting is quicker on [[brine|salt water]].

−

===Corrosion of coinage metals===

−

[[Coinage metal]]s, such as copper and silver, can also slowly corrode.

−

At standard temperature and pressure, a [[patina]] of green-blue [[copper carbonate]] forms on the surface of [[copper]]. [[Silver]] [[cutlery]] that is in contact with food can develop a layer of [[Silver sulfide]].

−

===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]]/[[cathode]] formation, if a [[electrical insulator|non-conducting]] material covers the metal contact with the [[electrolyte]] is not possible and corrosion will not occur.

−

====Coating====

−

Metals are [[coat|coated]] on its surface with [[paint]] or some other non-conducting coating. This prevents the [[electrolyte]] from reaching the metal surface '''IF''' the coating is complete. [[Scratch|Scratches]] exposing the metal will corrode with the region under the paint, adjacent to the scratch, to be [[anode|anodic]].

−

Other prevention is called ''[[passivation]]'' where a metal is coated with another metal such as [[tin can]]. Tin is a metal that rapidly corrodes to form a mono-molecular [[oxide]] coating 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 the can corrodes rapidly.

−

====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. 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.

−

[[Zinc]] bars are attached at various locations on steel [[ship]] [[Hull (watercraft)|hulls]] to render the ship hull [[cathode|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, buried or exposed an ingot of magnesium (or zinc) is [[bury|buried]] beside the [[Pipe (material)|pipeline]] and [[wire|connected electrically]] to the pipe above ground. The pipeline is forced to be a cathode and is protected. The magnesium anode is sacrificed. At intervals new [[ingot]]s are buried to replace those lost.

−

== See also ==

−

* [[Activity series of metals]]

−

* [[Bioelectricity]]

−

* [[Contact tension]] - a historical forerunner to the theory of electrochemistry.

−

* [[Electrochemical potential]]

−

* [[Frost diagram]]

−

* [[List of important publications in chemistry#Electrochemistry|Important publications in electrochemistry]]

−

* [[Pourbaix diagram]]

−

* [[Redox titration]]

−

* [[Table of standard electrode potentials]]

−

==

−

== External links ==

−

* [http://www.electrochemistry.net Electrochemistry.net]

−

* [http://www.electrochem.org ECS (The Electrochemical Society)]

−

* [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://www.abc.chemistry.bsu.by/vi/ Potentiodynamic Electrochemical Impedance Spectroscopy]

−

* [http://www.nanoelectrode.com Nanoelectrode.com] News and research articles related to nanoelectrochemistry

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[[Category:Electrochemistry|Electrochemistry]]

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[[Category:Physical chemistry]]

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Changes: Electrochemistry

Latest revision as of 10:24, 20 June 2013

@@ Line 1: / Line 1: @@
+#redirect[[Bioelectrochemistry]]
-{{Biopsy}}
-[[Image:Faraday-Daniell.PNG|thumb|250px|English chemists [[John Frederic Daniell|John Daniell]] ([[relative direction|left]]) and [[Michael Faraday]] ([[relative direction|right]]), both credited to be founders of electrochemistry as known today.]]
-'''Electrochemistry''' is a branch of [[chemistry]] that studies the reactions which take place at the interface of an electronic [[Electrical conductor|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 (electricity)|battery]], it is an ''electrochemical'' reaction. In general, electrochemistry deals with situations where an [[oxidation]] and a [[redox|reduction]] reaction are separated in space. The direct [[charge transfer]] from one molecule to another is not the topic of electrochemistry.
-==History==
-{{main|History of electrochemistry}}
-===16th to 18th century developments===
-[[Image:Guericke-electricaldevice.PNG|thumb|200px|left|[[Germany|German]] [[physicist]] [[Otto von Guericke]] beside his electrical generator while conducting an experiment.]]
-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 [[Germany|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 electricity|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 [[France|French]] [[chemist]] [[C.F. du Fay|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"'' (from the [[Latin language|Latin]] for ''"glass"''), or positive, electricity; and ''"resinous,"'' or negative, electricity. This was the ''two-fluid theory'' of electricity, which was to be opposed by [[Benjamin Franklin|Benjamin Franklin's]] ''one-fluid theory'' later in the century.[[Image:Galvani-frog-legs.PNG|thumb|left|200px|Late [[1780s]] diagram of Galvani's experiment on frog legs.]]
-[[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.
-[[Image:Volta-and-napoleon.PNG|thumb|right|200px|[[Italy|Italian]] [[physicist]] [[Alessandro Volta]] showing his ''"[[Battery (electricity)|battery]]"'' to [[France|French]] [[emperor]] [[Napoleon I of France|Napoleon Bonaparte]] in early [[1800s]].]]
-In the late [[1700s]] the [[Italy|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 [[nerve]]s and [[muscle]]s  spanned by [[metal]] [[probe]]s. 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 [[Electric ray|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.
-===19th century===
-[[Image:Humphrydavy.jpg|thumb|left|150px|Sir Humphry Davy's portrait in [[1800s]].]]
-In [[1800]], the English chemists [[William Nicholson (chemist)]] 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 [[thermoelectricity|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 cell]]s 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.
-[[Image:ChemicalHistoryofaCandle.PNG|thumb|right|140px|Professor Michael Faraday's portrait on his book [[The Chemical History of a Candle]].]]
-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 [[Ohm's law|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. In [[1836]] [[John Frederic Daniell|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 [[alloy]]ing the [[amalgam|amalgamated]] [[zinc]] of Sturgeon with [[Mercury (element)|mercury]] would produce a better voltage.
-[[Image:Arrhenius2.jpg|thumb|left|140px|Swedish chemist [[Svante Arrhenius]] portrait circa [[1880s]].]]
-[[William Robert Grove|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 battery|zinc carbon cell]].
-[[Svante August Arrhenius]] published 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 [[electrolyte]]s, when dissolved in water, become to varying degrees split or dissociated into electrically opposite positive and negative ions.
-In [[1886]] [[Paul Héroult]] and [[Charles Martin Hall|Charles M. Hall]] developed a successful method to obtain [[aluminum]] by using the principles described by Michael Faraday.
-In [[1894]] [[Wilhelm Ostwald|Friedrich Ostwald]] concluded important studies of the [[electrical conductivity]] and electrolytic dissociation of [[organic acid]]s.
-[[Image:Walther Nernst 2.jpg|thumb|right|140px|German scientist [[Walther Nernst]] portrait in [[1910s]].]]
-[[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 [[1902]], [[The Electrochemical Society]] (ECS) was founded.
-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]] (ISE) was founded.
-By the [[1960s]]&ndash;[[1970s]] [[quantum electrochemistry]] was developed by [[Revaz Dogonadze]] and his pupils.
-==Electrolysis==
-{{Main|Electrolysis}}
-Spontaneous redox reactions produces electricity, thus passage of electrons through a wire in the [[electric circuit]]. Electrolysis requires an external source of [[electrical energy]] to induce a chemical reaction, this process takes place in a compartment called [[electrolytic cell]]. Principles involved on electrolysis are the same as featured on electrochemical cells.
-===Electrolysis of molten sodium chloride===
-<!-- Image with unknown copyright status removed: [[Image:Downs sodium productioncell.jpg|thumb|250px|right|Down's cell diagram.]] -->
-When molten, [[sodium chloride]] can be electrolysed to yield metallic [[sodium]] and gaseous [[chlorine]]. Industrially this process takes place in a special cell named Down's cell. The cell is connected to a battery, allowing [[electrons]] [[to migrate|migration]] from the battery to the electrolytic cell.
-Reactions that take place at Down's cell are the following:
-:<math>\mbox{Anode (oxidation): }2Cl^{-} \rightarrow Cl_{2}(g) + 2e^{-}\,</math>
-:<math>\mbox{Cathode (reduction): }2Na^{+}(l) + 2e^{-} \rightarrow 2Na(l)\,</math>
-:<math>\mbox{Overall reaction: }2Na^{+} + 2Cl^{-}(l) \rightarrow 2Na(l) + Cl_{2}(g)\,</math>
-This process can yield industrial amounts of metallic sodium and gaseous chlorine, and is widely used on [[mineral dressing]] and [[metallurgy]] [[industry|industries]].
-Standard [[emf]] for this process is approximately -4 [[V]] indicating a non-spontaneous process. In order this reaction to occur the battery should provide at least a potential of 4V. However, on mineral refining industry, higher voltages are used, due to low efficiency on the process.
-===Electrolysis of water===
-[[Image:Hoffman voltameter.jpg|thumb|190px|Diagram of a Hofmann voltameter, showing electrolysis of water.]]
-{{Main|Electrolysis of water}}
-[[Water]] at standard temperature and pressure conditions doesn't decompose into [[hydrogen]] and [[oxygen]] [[Spontaneous process|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]].  In it, a pair of inert [[electrodes]] usually made of [[platinum]] act as anode and cathode in the electrolytic process. After the water (if pure) has been placed in the [[apparatus]], nothing happens, hence there are not enough [[ions]] to let the passage of electrons occur. To start the electrolysis an electrolyte should be placed in, usually [[sodium chloride]] or [[sulfuric acid]] (most used 0.1 [[Molar concentration|M]]).
-Bubbles from the gases will be seen near both electrodes. The following half reactions describe the process mentioned above:
-:<math>\mbox{Anode (oxidation): }2H_{2}O(l) \rightarrow O_{2}(g) + 4H^{+}(aq) + 4e^{-}\,</math>
-:<math>\mbox{Cathode (reduction): }2H_{2}O(g) + 2e^{-} \rightarrow H_{2}(g) + 2OH^{-}(aq)\,</math>
-:<math>\mbox{Overall reaction: }2H_{2}O(l) \rightarrow 2H_{2}(g) + O_{2}(g)\,</math>
-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.  However, it is considered to be a complex process because the contents in solution have to be analyzed in  [[chemical reaction|half reactions]], whether reduced or oxidized.
-====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. Usually, water is electrolysed as mentioned in electrolysis of water yielding ''gaseous [[oxygen]] in the anode'' and gaseous [[hydrogen]] in the cathode. On the other hand, sodium chloride in water [[Dissociation (chemistry)|dissociates]] in Na<sup>+</sup> and Cl<sup>-</sup> ions, [[anion]] will be attracted to the cathode, thus reducing the [[sodium]] ion.  The [[cation]] will then be attracted to the anode oxidizing [[chloride]] ion.
-The following half reactions describes the process mentioned:
-:<math>\mbox{1. Cathode: }Na^{+}(aq)+ 1e^{-} \rightarrow Na(s) \qquad E^{o}_{red}=-2.71 V\,</math>
-:<math>\mbox{2. Anode: }2Cl^{-}(aq) \rightarrow Cl_{2}(g) + 2e^{-} \qquad E^{o}_{red}= +1.36 V\,</math>
-:<math>\mbox{3. Cathode: }2H_{2}O(l) + 2e^{+} \rightarrow H_{2}(g) + 2OH^{-}(aq)\qquad E^{o}_{red}=-0.83 V\,</math>
-:<math>\mbox{4. Anode: } 2H_{2}O(l) \rightarrow O_{2}(g) + 4H^{+}(aq) + 4e^{-}\qquad E^{o}_{red}=+1.23V\,</math>
-Reaction 1 is discarded as it has the most [[Negative and non-negative numbers|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<sup>-</sup> ion is favored for [[redox|reduction]], then the water reaction is favored for [[oxidation]] producing 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|overvoltage effect]]. Additional voltage is sometimes required, beyond the voltage predicted by the <math>E^{o}_{cell}\,</math>. This may be due to [[chemical kinetics|kinetic]] rather than [[Thermochemistry|thermodynamic]] considerations. In fact, it has been proven that the [[activation energy]] for the chloride ion is very low, hence favorable in [[chemical kinetics|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 [[voltage]] of the external source has to be increased (hence, overvoltage).
-Finally, reaction 3 is favorable because it describes the proliferation of [[Hydroxide|OH<sup>-</sup>]] ions thus letting a probable reduction of [[Proton|H<sup>+</sup>]] ions less favorable an option.
-The overall reaction for the process according to the analysis would be the following:
-:<math>\mbox{Anode (Oxidation): } 2Cl^{-}(aq)\rightarrow Cl_{2}(g) + 2e^{-}\,</math>
-:<math>\mbox{Cathode (Reduction): } 2H_{2}O(l) + 2e{-}\rightarrow H_{2}(g) + 2OH^{-}(aq)\,</math>
-:<math>\mbox{Overall reaction: } 2H_{2}O + 2Cl^{-}(aq) \rightarrow H_{2}(g) + Cl_{2}(g) + 2OH^{-}(aq)\,</math>
-As the overall reaction indicates, the [[concentration]] of chloride ions is reduced in comparison to OH<sup>-</sup> ions (whose concentration increases).  The reaction also shows the production of gaseous [[hydrogen]], [[chlorine]] and aqueous [[sodium hydroxide]].
-===Quantitative electrolysis & Faraday Laws===
-{{Main|Faraday's law of electrolysis}}
-Quantitative aspects of electrolysis were originally developed by [[Michael Faraday]] in [[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]].
-[[Image:Nickel-electroplating.PNG|thumb|200px|right|To [[refining|refine]] metals a process named [[electroplating]] is used (diagram shows [[Nickel]] refining); the process has its bases on the first and the second law of electrolysis stated by Faraday in the [[19th century]].]]
-====First law====
-Faraday concluded after several experiments on [[electrical current]] in [[spontaneous process|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:
-:<math>m \ = \ { 1 \over F \ } \cdot { Q M \over n } </math>
-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 [[coulomb]]s),
-:''n'' is the valence number of the substance as an ion in solution (electrons per ion),
-:''F'' is the Faraday's constant = 96,485 [[coulomb]]s per mole,
-:''M'' is the molar mass of the substance (in grams per [[mole (unit)|mole]]).
-Faraday's constant is equal to the amount of charge carried by one mole of electrons.
-====Second law====
-{{Main|Electroplating}}
-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 [[ratio]] of their chemical [[equivalent weight]]s.
-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 [[metal]]s to avoid corrosion.
-==Principles==
-===Redox reactions===
-{{main|Redox reaction}}
-Electrochemical process are redox reactions where [[energy]] is produced by a [[Spontaneous process|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 [[chemical element|element]]s involved in an electrochemical [[chemical reaction|reaction]] are characterized by the number of [[electron]]s each has. The ''oxidation state'' of an [[ion]] is the number of electrons it has accepted or donated compared to its neutral state (which is defined as having an oxidation state of 0). 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 (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 of a substance is called [[oxidation]], and the gain of electrons is [[redox|reduction]]. This can be easily remembered through the use of [[mnemonic]] devices. 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).
-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 (it has a higher [[electronegativity]]) than oxygen.
-===Balancing redox reactions===
-{{main|Chemical equation}}
-Electrochemical reactions in water are better understood by balancing redox reactions using the [[Ion-Electron Method]] where [[Proton|H<sup>+</sup>]] , [[Hydroxide|OH<sup>-</sup>]] ion, [[Water (molecule)|H<sub>2</sub>O]] and electrons (to compensate the oxidation changes) are added to cell's [[half reaction]]s for oxidation and reduction.
-====Acid medium====
-In acid medium [[Proton|H]] atoms and water are added to [[half reaction]]s to balance the overall reaction.
-For example, when [[Manganese]] reacts with [[Sodium bismuthate]].
-:<math>\mbox{Reaction unbalanced: }\mbox{Mn}^{2+}(aq) + \mbox{NaBiO}_3(s)\rightarrow\mbox{Bi}^{3+}(aq) + \mbox{MnO}_4^{-}(aq)\,</math>
-:<math>\mbox{Oxidation: }\mbox{4H}_2\mbox{O}(l)+\mbox{Mn}^{2+}(aq)\rightarrow\mbox{MnO}_4^{-}(aq) + \mbox{8H}^{+}(aq)+\mbox{5e}^{-}\,</math>
-:<math>\mbox{Reduction: }\mbox{2e}^{-}+ \mbox{6H}^{+}(aq) + \mbox{BiO}_3^{-}(s)\rightarrow\mbox{Bi}^{3+}(aq) + \mbox{3H}_2\mbox{O}(l)\,</math>
-Finally the reaction is balanced by [[multiplication|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.
-:<math>\mbox{8H}_2\mbox{O}(l)+\mbox{2Mn}^{2+}(aq)\rightarrow\mbox{2MnO}_4^{-}(aq) + \mbox{16H}^{+}(aq)+\mbox{10e}^{-}\,</math>
-:<math>\mbox{10e}^{-}+ \mbox{30H}^{+}(aq) + \mbox{5BiO}_3^{-}(s)\rightarrow\mbox{5Bi}^{3+}(aq) + \mbox{15H}_2\mbox{O}(l)\,</math>
-Reaction balanced:
-:<math>\mbox{14H}^{+}(aq) + \mbox{2Mn}^{2+}(aq)+ \mbox{5NaBiO}_3(s)\rightarrow\mbox{7H}_2\mbox{O}(l) + \mbox{2MnO}_4^{-}(aq)+\mbox{5Bi}^{3+}(aq)+\mbox{5Na}^{+}(aq)\,</math>
-====Basic medium====
-In basic medium [[Hydroxide|OH<sup>-</sup>]] ions and [[Water (molecule)|water]] are added to half reactions to balance the overall reaction. For example on reaction between [[Potassium permanganate]] and [[Sodium sulfite]].
-:<math>\mbox{Reaction unbalanced: }\mbox{KMnO}_{4}+\mbox{Na}_{2}\mbox{SO}_3+\mbox{H}_2\mbox{O}\rightarrow\mbox{MnO}_{2}+\mbox{Na}_{2}\mbox{SO}_{4}+\mbox{KOH}\,</math>
-:<math>\mbox{Reduction: }\mbox{3e}^{-}+\mbox{2H}_{2}\mbox{O}+\mbox{MnO}_{4}^{-}\rightarrow\mbox{MnO}_{2}+\mbox{4OH}^{-}\,</math>
-:<math>\mbox{Oxidation: }\mbox{2OH}^{-}+\mbox{SO}^{2-}_{3}\rightarrow\mbox{SO}^{2-}_{4}+\mbox{H}_{2}\mbox{O}+\mbox{2e}^{-}\,</math>
-The same procedure as followed on acid medium by multiplying electrons to opposite half reactions solve the equation thus balancing the overall reaction.
-:<math>\mbox{6e}^{-}+\mbox{4H}_{2}\mbox{O}+\mbox{2MnO}_{4}^{-}\rightarrow\mbox{2MnO}_{2}+\mbox{8OH}^{-}\,</math>
-:<math>\mbox{6OH}^{-}+\mbox{3SO}^{2-}_{3}\rightarrow\mbox{3SO}^{2-}_{4}+\mbox{3H}_{2}\mbox{O}+\mbox{6e}^{-}\,</math>
-Equation balanced:
-:<math>\mbox{2KMnO}_{4}+\mbox{3Na}_{2}\mbox{SO}_3+\mbox{H}_2\mbox{O}\rightarrow\mbox{2MnO}_{2}+\mbox{3Na}_{2}\mbox{SO}_{4}+\mbox{2KOH}\,</math>
-====Neutral medium====
-The same procedure as used on acid medium is applied, for example on balancing using electron ion method to [[Combustion|complete combustion]] of [[propane gas]].
-:<math>\mbox{Reaction unbalanced: }\mbox{C}_{3}\mbox{H}_{8}+\mbox{O}_{2}\rightarrow\mbox{CO}_{2}+\mbox{H}_{2}\mbox{O}\,</math>
-:<math>\mbox{Reduction: }\mbox{4H}^{+} + \mbox{O}_{2}\rightarrow\mbox{H}_{2}\mbox{O}+\mbox{H}_{2}\mbox{O}+ \mbox{4e}^{-}\,</math>
-:<math>\mbox{Oxidation: }\mbox{20e}^{-}+\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\mbox{20H}^{+}\,</math>
-As in acid and basic medium, electrons which were used to compensate oxidation changes are multiplied to opposite half reactions, thus solving the equation.
-:<math>\mbox{20H}^{+}+\mbox{5O}_{2}\rightarrow\mbox{5H}_{2}\mbox{O}+\mbox{5H}_{2}\mbox{O}+\mbox{20e}^{-}\,</math>
-:<math>\mbox{20e}^{-}+\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\mbox{20H}^{+}\,</math>
-Equation balanced:
-:<math>\mbox{C}_{3}\mbox{H}_{8}+\mbox{5O}_{2}\rightarrow\mbox{3CO}_{2}+\mbox{4H}_{2}\mbox{O}\,</math>
-==Electrochemical cells==
-{{main|Electrochemical cell}}
-[[Image:Galvanic cell.png|thumb|right|260px|A modified version of [[Galvanic cell|Daniells Cell]]s, a U—Shaped tube is replaced with a porous disk acting as [[saline bridge]] thus electric current is produced.]]
-An electrochemical cell is a device capable of producing electric current from energy released by a [[Spontaneous process|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 who 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|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.
-An electrochemical cell which [[electrode]]s are [[Zinc]] and [[Copper]] submerged on [[Zinc sulfate]] and [[Copper sulfate]] respectively is known as [[Daniell cell|Daniells cell]].
-Half reactions for a Daniells cell are these:
-:<math>\mbox{Zinc electrode (anode) : }\mbox{Zn}(s)\rightarrow\mbox{Zn}^{2+}(aq)+\mbox{2e}^{-}\,</math>
-:<math>\mbox{Copper electrode (cathode) : }\mbox{Cu}^{2+}(aq)+\mbox{2e}^{-}\rightarrow\mbox{Cu}(s)\,</math>
-[[Image:BASi epsilon C3 cell stand.jpg|thumb|right||260px|A modern cell stand for electrochemical research.  The electrodes attach to high-quality metallic wires, and the stand is attached to a potentiostat/galvanostat (not pictured).  A [[shotglass]]-shaped container is [[Aerated water|aerated]] with a noble gas and sealed with the [[teflon]] block.]]
-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 [[Direct current|D.C.]] electric current.
-A [[galvanometer|voltameter]] is capable of measuring the change of [[Electric potential|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:
-:<math>\mbox{Zn}(s)|\mbox{Zn}^{2+}(1M)||\mbox{Cu}^{2+}(1M)|\mbox{Cu}(s)\,</math>
-First, the reduced form of the metal to be oxidized at the anode (Zn) is written . This is separated from its oxidised 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.
-==Standard electrode potential==
-{{Main|Standard electrode potential}}
-[[Image:Reduction-potentials2.PNG|thumb|right|260px|The [[standard reduction potentials]] table is determined in a modified version of [[galvanic cell]] using an [[Hydrogen]] [[electrode]] as [[cathode]], because Hydrogen is taken as reference, standard reduction potential for that substance is zero (gray [[highlighter|highlight]]).]]
-Standard electrode potential is the value of the standard [[emf]] of a cell in which molecular hydrogen under standard pressure (10<sup>5</sup> Pa) 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 [[redox|reduction]] takes place on the chosen electrode, hence standard electrode potential are [[tabulation|tabulated]] on reduction potentials, thus tables are built on [[standard reduction potential]]s noted as <math>\mbox{E}^{0}_{red}\,</math>.
-Standard cell potential is calculated by the difference between the standard reduction potentials of each electrode.
-:<math>\mbox{E}^{o}_{cell}=\mbox{E}^{o}_{red}(cathode)-\mbox{E}^{o}_{red}(anode)</math>
-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 <math>\mbox{E}^{0}_{red}=0\,</math>. Cell's half reaction used for this procedure is [[hydrogen]] which in [[Standard conditions for temperature and pressure|standard temperature and pressure]] conditions (10<sup>5</sup> Pa, 298.15&nbsp;K, 1 mol. L<sup>-1</sup>) acts as a zero volt electrode.
-The [[standard hydrogen electrode]] or ([[Standard hydrogen electrode|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]]. This setup is placed in a solution of [[Hydrochloric acid]], plenty of H<sup>+</sup> ions, gaseous [[hydrogen]] enter through the tube and react over the platinum blade thus allowing reduction and oxidation processes to occur.
-[[Standard hydrogen electrode|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:
-:<math>\mbox{Cell diagram}\,</math>
-:<math>\mbox{Pt}(s)|\mbox{H}_{2}(1 atm)|\mbox{H}^{+}(1 M)||\mbox{Cu}^{2+}(1 M)|\mbox{Cu}(s)\,</math>
-:<math>\mbox{E}^{o}_{cell}=\mbox{E}^{o}_{red}(cathode)-\mbox{E}^{o}_{red}(anode)</math>
-At standard temperature pressure conditions cell's [[electromotive force|emf]] (measured by a [[multimeter]]) is 0.34 V, conventionally [[Standard hydrogen electrode|SHE]] has a zero value, thus replacing on previous equation gives:
-:<math>\mbox{0.34V}_{cell}=\mbox{E}^{o}_{\mbox{Cu}^{2+}/\mbox{Cu}}-\mbox{E}^{o}_{\mbox{H}^{+}/\mbox{H}_{2}}</math>
-:<math>\mbox{0.34V}_{cell}=\mbox{E}^{o}_{\mbox{Cu}^{2+}/\mbox{Cu}}-0</math>
-Electrochemical cell's [[electromotive force|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 coefficient]]s on balanced cell equation will not change <math>\mbox{E}^{0}_{red}\,</math> value because standard electrode potential are [[Intensive and extensive properties|intensive properties]].
-==Spontaneity of Redox systems==
-{{main|Spontaneous process}}
-On electrochemical cells, [[chemical energy]] transforms into [[electrical energy]] and is expressed mathematically as the product between cell's emf by [[electrical charge]] in [[Coulomb]]s.
-:<math>\mbox{Electrical energy}=(\mbox{volts})(\mbox{coulombs})\,</math>
-:<math>\mbox{Electrical energy}=\mbox{joules}\,</math>
-Electrochemical cell's total charge is determined by multiplying the number of moles by [[Faraday's constant]] (F).
-:<math>\mbox{Total charge}=\mbox{n}\mbox{F}\,</math>
-Faraday's constant is the electrical charge in 1 [[Mole (unit)|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,
-:<math>\mbox{W}_{max}=\mbox{W}_{electrical}\,</math>
-:<math>\mbox{W}_{max}=-\mbox{nFE}_{cell}\,</math>
-,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 <math>\Delta G\,</math>gives the relation between spontaneity and electrochemical cells.
-:<math>\Delta G=-\mbox{nFE}_{cell}\,</math>
-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 (electricity)|current]], in [[electrochemical cell]]s. This is the basis of all [[battery (electricity)|batteries]] and [[fuel cell]]s. For example, gaseous oxygen (O<sub>2</sub>) and
-hydrogen (H<sub>2</sub>) can be combined in a fuel cell to form water and
-energy (a combination of heat and electrical energy, typically).
-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:
-:<math>\Delta G^{o}=\mbox{-RT ln K}\,</math>
-:<math>\mbox{-nFE}^{o}_{cell}=\mbox{-RT ln K}\,</math>
-Solving both equations express cell's mathematical relation between standard potential, and equilibrium constant.
-:<math>\mbox{E}^{o}_{cell}={\mbox{RT} \over \mbox{nF}} \mbox{ln K}\,</math>
-Previous equation can use [[Briggsian logarithm]] as shown below:
-:<math>\mbox{E}^{o}_{cell}={0.0592 \mbox{V} \over \mbox{n}} \mbox{log K}\,</math>
-==Cell emf dependency on changes in concentration==
-===Nernst Equation===
-{{Main|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.
-In the mid [[1800s]] [[Willard Gibbs]] formulated an equation for spontaneous process at any conditions,
-:<math>\Delta G=\Delta G^{o}+\mbox{RT ln Q}\,</math> ,
-Where:
-''ΔG'' = change in [[Gibbs free energy]], ''T'' = absolute [[temperature]], ''R'' = [[gas constant]], ln = [[natural logarithm]], ''Q'' = [[reaction quotient]].
-Willard stated Q's dependency over reactants and products activity and designated it as their respective [[Activity (chemistry)|chemical activity]].
-Walther based on Willard Gibbs work during the mid [[19th century]], formulated a new equation where replaced <math>\Delta G\,</math>'s value with cell's respective maximum electrical work, on Gibbs equation.
-:<math>nF\Delta E = nF\Delta E^\circ - R T \ln Q \, \,</math>
-Where:
-''n'' = number of [[electrons]]/[[Mole (unit)|mole]] product, ''F'' = [[Faraday constant]] ([[coulomb]]s/[[Mole (unit)|mole]]), and ''ΔE'' = [[electrical potential of the reaction]].
-Finally he replaced <math>-nF\Delta E\,</math>'s value with electrochemical cell potential, thus formulating a new equation which now bears his name.
-:<math>\Delta E=\Delta E^{o}- {\mbox{RT} \over \mbox{nF}} \mbox{ln Q}\,</math>
-Assuming standard conditions (<math>Temperature = 298 K , 25 C\,</math>) and [[Universal gas constant|R]] = <math>8.3145 {J \over K mol}</math> the equation above can be expressed on [[Common logarithm|Base—10 logarithm]] as shown below:
-:<math>\Delta E=\Delta E^{o}- {\mbox{0.0592 V} \over \mbox{n}} \mbox{log Q}\,</math>
-===Concentration cells===
-{{Main|Concentration cell}}
-[[Image:Cell-membrane-electrochemical.PNG|thumb|250px|Calculating [[membrane potential]] is good example where concentration cells are used in biology to understanding cell's [[metabolism]] such as [[Na-K pump|Na<sup>+</sup>(red) K<sup>+</sup>(blue) pump]].]]
-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|blue vitriol's]] solution, whose concentrations are 0.05 [[Molar concentration|M]] and 2.0 [[Molar concentration|M]] , while connected through wire and saline bridge.
-:<math>Cu^{2+}(aq)+2e^{-}\rightarrow \mbox{Cu}(s)</math>
-[[Le Chatelier's principle]] indicates reaction is favourable to reduction as concentration of <math>Cu^{2+}\,</math> 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:
-:<math>Cu(s)|Cu^{2+}(0.05 M)||Cu^{2+}(2.0 M)|Cu(s)\,</math>
-Where both half cell reactions for oxidation and reduction are:
-:<math>Oxidation: Cu(s)\rightarrow \mbox{Cu}^{2+} (0.05 M) + 2e^{-}\,</math>
-:<math>Reduction: Cu^{2+} (2.0 M) +2e^{-} \rightarrow \mbox{Cu} (s)\,</math>
-:<math>Overall reaction: Cu^{2+} (2.0 M) \rightarrow \mbox{Cu}^{2+} (0.05 M)\,</math>
-Where cell's emf is calculated through Nernst equation as follows:
-:<math>E = E^{o}- {0.0257 V \over 2} ln {[Cu^{2+}]_{diluted}\over [Cu^{2+}]_{concentrated}}\,</math>
-<math>E^{o}\,</math>'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:
-:<math>E = 0- {0.0257 V \over 2} ln {0.05\over 2.0}\,</math>
-:<math>E = 0.0474 V\,</math>
-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 [[Synapses|nerve synapses]] and [[Cardiac cycle|cardiac beat]].
-==Battery==
-{{Main|Battery (electricity)}}
-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===
-{{Main|Dry cell}}
-[[Image:Zincbattery.png|thumb|250px|Zinc carbon battery diagram.]]
-Dry cells don't have a [[fluid]] electrolyte instead they use a moist electrolyte paste. [[Zinc-carbon battery|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.
-[[Georges Leclanché|Leclanché's]] simplified half reactions are shown below:
-:<math>Anode: Zn(s) \rightarrow Zn^{2+} (aq) + 2e^{-}\,</math>
-:<math>Cathode: 2NH^{+}_{4}(aq)+ 2MnO_{2}(s) + 2e^{-}\rightarrow Mn_{2}O_{3}(s) + 2NH_{3} (aq) + H_{2}O (l)\,</math>
-:<math>\mbox{Overall reaction:}\,</math>
-:<math>Zn(s) + 2NH^{+}_{4}(aq)+ 2MnO_{2}(s) \rightarrow Zn^{2+}(aq) + Mn_{2}O_{3}(s) + 2NH_{3} (aq) + H_{2}O (l)\,</math>
-The voltage obtained from the [[zinc-carbon battery]] is 1.5 [[Volt|V]] approximately.
-===Mercury battery===
-{{Main|Mercury battery}}
-[[Image:Mercurybattery2.PNG|thumb|200px|Cutaway view of a Mercury battery diagram.]]
-Mercury battery has many applications on [[medicine]] and [[electronics]]. The battery consists of a [[steel]]—made container with the shape of a cylinder acting as the cathode, where an [[amalgam|amalgamated]] anode of mercury and zinc is surrounded by a stronger alkaline electrolyte and a paste of [[Zinc oxide]] and [[Mercury(II) oxide]] .
-Mercury battery half reactions are shown below:
-:<math>Anode: Zn(Hg) + 2OH^{-} (aq) \rightarrow ZnO(s) + H_{2}O (l) + 2e^{-}\,</math>
-:<math>Cathode: HgO(s) + H_{2}O(l) + 2e^{-}\rightarrow Hg(l) + 2OH^{-} (aq)\,</math>
-:<math>\mbox{Overall reaction:}\,</math>
-:<math>Zn(Hg) + HgO(s) \rightarrow ZnO(s) + Hg(l)\,</math>
-There are no changes on the electrolyte's composition when cell works. Mercurium battery provides 1.35 V of [[direct current]].
-===Lead-acid battery===
-[[Image:Lead acid cell.jpg|thumb|100px|A sealed Lead acid battery.]]
-{{Main|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:
-:<math>Anode: Pb(s) + SO^{2-}_{4}(aq) \rightarrow PbSO_{4}(s) + 2e^{-}\,</math>
-:<math>Cathode: PbO_{2}(s) + 4H^{+}(aq) + SO^{2-}_{4}(aq) + 2e^{-} \rightarrow PbSO_{4}(s) + 2H_{2}O(l)\,</math>
-<math>\mbox{Overall reaction:} Pb(s) + PbO_{2}(s) + 4H^{+}(aq)+2SO^{2-}_{4}(aq) \rightarrow 2PbSO_{4}(s) + 2H_{2}O(l)</math>
-At standard conditions, each cell may produce a [[direct current]] of 2 [[Volts|V]], hence overall voltage produced is 12 V. Lead-acid batteries, differing from Mercury and Zinc-carbon batteries, are [[Rechargeable battery|rechargeable]]. If an external voltage is supplied to the battery it will produce an [[electrolysis]] of the products in the overall reaction (discharge), thus recovering initial components which made the battery work.
-===Solid state Lithium battery===
-{{Main|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 [[faraday constant|mole of electrons]]. This battery is rechargeable and it can provide a [[direct current]] of about 3 [[Volts|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===
-{{Main|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 battery]] and [[zinc bromine battery]].
-===Fuel cells===
-{{Main|Fuel cell}}
-[[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[[percentage|%]] of the original [[chemical energy]] when [[combustion|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 electrode|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:
-:<math>Anode: 2H_{2}(g)+ 4OH^{-}(aq)\rightarrow 4H_{2}O(l)+4e^{-}\,</math>
-:<math>Cathode: O_{2}(g)+ 2H_{2}O(l) + 4e^{-}\rightarrow 4OH^{-}(aq)\,</math>
-:<math>\mbox{Overall reaction:} 2H_{2}(g) + O_{2}(g)\rightarrow 2H_{2}O(l)\,</math>
-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 [[electrocatalyst]]s. [[Platinum]], [[nickel]], [[rhodium]] are good electrocatalysts.
-==Corrosion==
-{{Main|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 of reddish rust.  Other examples include the black tarnish on [[silver]], and red or green corrosion that may appear on [[copper]] and its alloys, such as [[brass]]. The cost of replacing metals lost to corrosion is in the multi-billions of [[American dollar|dollars]] per year.
-===Iron corrosion===
-[[Image:Iron-rusting-scheme.PNG|thumb|270px|right|Diagram showing a water [[drop|droplet]] over an iron surface. Electrochemical mechanisms involved develop iron rusting process.{{replacethisimage}}]]
-For iron rust to occur the metal has to be in contact with [[oxygen]] and [[water]], although [[chemical reaction]]s 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.
-###:<math>Fe(s)\rightarrow Fe^{2+}(aq) + 2e^{-}\,</math>
-##[[Electrons]] are transferred from [[iron]] reducing oxygen in the [[atmosphere]] into [[water (molecule)|water]] on the cathode, which is placed in another region of the metal.
-###:<math>O_{2}(g) + 4H^{+}(aq) + 4e^{-} \rightarrow 2H_{2}O(l)\,</math>
-##Global reaction for the process:
-##:<math>2Fe(s) + O_{2}(g) + 4H^{+}(aq) \rightarrow 2Fe^{2+}(aq) + 2H_{2}O(l)\,</math>
-##Standard [[emf]] for iron rusting:
-###:<math>E^{o}=E^{o}_{cathode}-E^{o}_{anode}\,</math>
-###:<math>E^{o}=1.23V-(-0.44V)=1.67V\,</math>
-Iron corrosion takes place on acid medium; [[Proton|H<sup>+</sup>]] [[ions]] come from reaction between [[carbon dioxide]] in the atmosphere and water, forming [[carbonic acid]]. Fe<sup>2+</sup> ions oxides, following this equation:
-:<math>4Fe^{2+}(aq) + O_{2}(g) + (4+2x)H_{2}O(l) \rightarrow 2Fe_{2}O_{3}.xH_{2}O + 8H^{+}(aq)</math>
-[[Iron(III) oxide]] [[hydrated]] is known as rust. Water associated with iron oxide it varies, thus chemical representation is presented as <math>Fe_{2}O_{3}.xH_{2}O\,</math>.
-The [[electric circuit]] works as passage of electrons and ions occurs, thus if an electrolyte is present it will facilitate [[oxidation]], this explains why rusting is quicker on [[brine|salt water]].
-===Corrosion of coinage metals===
-[[Coinage metal]]s, such as copper and silver, can also slowly corrode.
-At standard temperature and pressure, a [[patina]] of green-blue [[copper carbonate]] forms on the surface of [[copper]]. [[Silver]] [[cutlery]] that is in contact with food can develop a layer of [[Silver sulfide]].
-===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]]/[[cathode]] formation, if a [[electrical insulator|non-conducting]] material covers the metal contact with the [[electrolyte]] is not possible and corrosion will not occur.
-====Coating====
-Metals are [[coat|coated]] on its surface with [[paint]] or some other non-conducting coating. This prevents the [[electrolyte]] from reaching the metal surface '''IF''' the coating is complete. [[Scratch|Scratches]] exposing the metal will corrode with the region under the paint, adjacent to the scratch, to be [[anode|anodic]].
-Other prevention is called ''[[passivation]]'' where a metal is coated with another metal such as [[tin can]]. Tin is a metal that rapidly corrodes to form a mono-molecular [[oxide]] coating 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 the can corrodes rapidly.
-====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. 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.
-[[Zinc]] bars are attached at various locations on steel [[ship]] [[Hull (watercraft)|hulls]] to render the ship hull [[cathode|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, buried or exposed an ingot of magnesium (or zinc) is [[bury|buried]] beside the [[Pipe (material)|pipeline]] and [[wire|connected electrically]] to the pipe above ground. The pipeline is forced to be a cathode and is protected. The magnesium anode is sacrificed. At intervals new [[ingot]]s are buried to replace those lost.
-== See also ==
-* [[Activity series of metals]]
-* [[Bioelectricity]]
-* [[Contact tension]] - a historical forerunner to the theory of electrochemistry.
-* [[Electrochemical potential]]
-* [[Frost diagram]]
-* [[List of important publications in chemistry#Electrochemistry|Important publications in electrochemistry]]
-* [[Pourbaix diagram]]
-* [[Redox titration]]
-* [[Table of standard electrode potentials]]
-==
-== External links ==
-* [http://www.electrochemistry.net Electrochemistry.net]
-* [http://www.electrochem.org ECS (The Electrochemical Society)]
-* [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://www.abc.chemistry.bsu.by/vi/ Potentiodynamic Electrochemical Impedance Spectroscopy]
-* [http://www.nanoelectrode.com Nanoelectrode.com] News and research articles related to nanoelectrochemistry
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