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| − | [[Image:Faraday-Daniell.PNG|thumb|English chemists [[John Frederic Daniell|John Daniell]] ([[relative direction|left]]) and [[Michael Faraday]] (right), both credited as founders of electrochemistry today.]]
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| − | '''Electrochemistry''' is a branch of [[chemistry]] involving the study of interrelationships between [[electricity]] and [[chemical reaction]]s. The chemical reactions generally take place in [[solution]], at the interface between an electron [[Electrical conductor|conductor]] (a [[metal]] or [[semiconductor]]) and an ion conductor (the [[electrolyte]]), and involve electron transfer between the electrode and the electrolyte or species in solution.
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| − | If a chemical reaction is driven by an external applied [[voltage]], as in [[electrolysis]], or if a voltage is generated by a chemical reaction, as in a [[battery (electricity)|battery]], the reaction is called an '''electrochemical reaction'''. Chemical reactions where electrons are transferred between [[molecule]]s are called oxidation/reduction ([[redox]]) reactions.
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| − | Electrochemical reactions are valuable for many important applications. For example, they may be used to extract [[metal]]s from their [[ore]]s, or to coat objects with metals or metal oxides through electrodeposition. The redox reaction may be used to detect alcohol in drunken drivers or to measure [[glucose]] levels in the blood of diabetics. In nature, the generation of chemical energy through [[photosynthesis]] is an electrochemical process.
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| − | ==History==
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| − | {{main|History of electrochemistry}}
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| − | === Developments from the sixteenth to eighteenth centuries ===
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| − | [[Image:Guericke-electricaldevice.PNG|thumb|left|[[Germany|German]] [[physicist]] [[Otto von Guericke]] beside his electrical generator while conducting an experiment.]]
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| − | The sixteenth century marked the beginning of electrical understanding. During that century 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—eighteenth century, 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. DuFay 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|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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| − | [[Image:Volta-and-napoleon.PNG|thumb|right|[[Italy|Italian]] [[physicist]] [[Alessandro Volta]] showing his ''"[[Battery (electricity)|battery]]"'' to [[France|French]] [[emperor]] [[Napoleon I of France|Napoleon Bonaparte]] in the early nineteenth century.]]
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| − | In the late eighteenth century, 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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| − | In 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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| − | ===Nineteenth century===
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| − | [[Image:Humphrydavy.jpg|thumb|left|upright|Sir Humphry Davy's portrait in the nineteenth century.]]
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| − | In 1800, [[William Nicholson (chemist)|William Nicholson]] and [[Johann Wilhelm Ritter]] succeeded in decomposing water into [[hydrogen]] and [[oxygen]] by [[electrolysis]]. Soon thereafter Ritter discovered the process of [[electroplating]]. He also observed that the amount of metal deposited and the amount of oxygen produced during an electrolytic process depended on the distance between the [[electrodes]]. By 1801 Ritter observed [[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]]. 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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| − | 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 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]]ated [[zinc]] of Sturgeon with [[Mercury (element)|mercury]] would produce a better voltage.
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| − | [[Image:Arrhenius2.jpg|thumb|left|upright|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 [[aluminium]] 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|upright|German scientist [[Walther Nernst]] portrait in the 1910s.]]
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| − | [[Hermann Nernst|Walther 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 [[Thermodynamic free energy|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. He also 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 twentieth century ===
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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 and some years later he was awarded 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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| − | ==Principles==
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| − | ===Redox reactions===
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| − | {{main|Redox reaction}}
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| − | Electrochemical processes involve [[redox]] reactions where an [[electron]] is transferred to or from a [[molecule]] or [[ion]] changing its [[oxidation state]]. This reaction can occur through the application of an external [[voltage]] or through the release of chemical energy.
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| − | ===Oxidation and reduction===
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| − | The atoms, ions, or molecules involved in an electrochemical [[chemical reaction|reaction]] are characterized by the number of [[electron]]s each has compared to its number of [[proton]]s called its '' oxidation state'' and is denoted by a + or a -. Thus the superoxide ion, O<sub>2</sub><sup>-</sup>, has an ''oxidation state'' of -1. An atom or ion that gives up an electron to another atom or ion has its oxidation state increase, and the recipient of the negatively charged electron has its oxidation state decrease. Oxidation and reduction always occur in a paired fashion such that one species is oxidized when another is reduced. This paired electron transfer is called a [[redox]] reaction.
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| − | For example when atomic [[sodium]] reacts with atomic [[chlorine]], sodium donates one electron and attains an oxidation state of +1. Chlorine accepts the electron and its oxidation state is reduced to −1. The sign of the oxidation state (positive/negative) actually corresponds to the value of each ion's electronic charge. The attraction of the differently charged sodium and chlorine ions is the reason they then form an [[ionic bond]].
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| − | The loss of electrons from an atom or molecule 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). For cases where electrons are shared (covalent bonds) between atoms with large differences in [[electronegativity]], the electron is assigned to the atom with the largest electronegativity in determining the oxidation state.
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| − | The atom or molecule which loses electrons is known as the ''reducing agent'', or ''reductant'', and the substance which accepts the electrons is called the ''oxidizing agent'', or ''oxidant''. The oxidizing agent is always being reduced in a reaction; the reducing agent is always being oxidized. Oxygen is a common oxidizing agent, but not the only one. Despite the name, an oxidation reaction does not necessarily need to involve oxygen. In fact, a [[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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| − | For reactions involving oxygen, the gain of oxygen implies the oxidation of the atom or molecule to which the oxygen is added (and the oxygen is reduced). For example, in the oxidation of [[octane]] by [[oxygen]] to form [[carbon dioxide]] and [[water]], both the carbon in the octane and the oxygen begin with an oxidation state of 0. In forming CO<sub>2</sub> the carbon loses four electrons to become C<sup>4+</sup> and the oxygens each gain two electrons to be O<sup>2-</sup>. In organic compounds, such as [[butane]] or [[ethanol]], the loss of hydrogen implies oxidation of the molecule from which it is lost (and the hydrogen is reduced). This follows because the hydrogen donates its electron in covalent bonds with non-metals but it takes the electron along when it is lost. Conversely, loss of oxygen or gain of hydrogen implies reduction.
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| − | ===Balancing redox reactions===
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| − | {{main|Chemical equation}}
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| − | Electrochemical reactions in water are better understood by balancing redox reactions using the [[Ion-Electron Method]] where [[hydronium|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.
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| − | ====Acid medium====
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| − | In acid medium [[hydronium|H+]] ions 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<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]].
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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]].
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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}+ \mbox{4e}^{-}\rightarrow\mbox{2H}_{2}\mbox{O}\,</math>
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| − | :<math>\mbox{Oxidation: }\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\mbox{20e}^{-}+\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}+\mbox{20e}^{-}\rightarrow\mbox{10H}_{2}\mbox{O}\,</math>
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| − | :<math>\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\mbox{20e}^{-}+\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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| − | {{main|Electrochemical cell}}
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| − | An electrochemical cell is a device that produces an electric current from energy released by a [[Spontaneous process|spontaneous]] redox reaction. This kind of cell includes the [[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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| − | Electrochemical cells have two conductive electrodes (the anode and the cathode). The [[anode]] is defined as the electrode where oxidation occurs and the [[cathode]] is the electrode where the reduction takes place. Electrodes can be made from any sufficiently conductive materials, such as metals, semiconductors, graphite, and even [[conductive polymer]]s. In between these electrodes is the [[electrolyte]], which contains ions that can freely move.
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| − | The Galvanic cell uses two different metal electrodes, each in an electrolyte where the positively charged ions are the oxidized form of the electrode metal. One electrode will undergo oxidation (the anode) and the other will undergo reduction (the cathode). The metal of the anode will oxidize, going from an oxidation state of 0 (in the solid form) to a positive oxidation state and become an ion. At the cathode, the metal ion in solution will accept one or more electrons from the cathode and the ion's oxidation state is reduced to 0. This forms a solid metal that [[electroplating|electrodeposits]] on the cathode. The two electrodes must be electrically connected to each other, allowing for a flow of electrons that leave the metal of the anode and flow through this connection to the ions at the surface of the cathode. This flow of electrons is an electrical current that can be used to do work, such as turn a motor or power a light.
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| − | A Galvanic cell whose [[electrode]]s are [[zinc]] and [[copper]] submerged in [[zinc sulfate]] and [[copper sulfate]], respectively, is known as a [[Daniell cell]].
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| − | Half reactions for a Daniell 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 [[Polytetrafluoroethylene|Teflon]] block.]]
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| − | In this example, the anode is zinc metal which oxidizes (loses electrons) to form zinc ions in solution, and copper ions accept electrons from the copper metal electrode and the ions deposit at the copper cathode as an electrodeposit. This cell forms a simple battery as it will spontaneously generate a flow of electrical current from the anode to the cathode through the external connection. This reaction can be driven in reverse by applying a voltage, resulting in the deposition of zinc metal at the anode and formation of copper ions at the cathode.
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| − | To provide a complete electric circuit, there must also be an ionic conduction path between the anode and cathode electrolytes in addition to the electron conduction path. The simplest ionic conduction path is to provide a liquid junction. To avoid mixing between the two electrolytes, the liquid junction can be provided through a porous plug that allows ion flow while reducing electrolyte mixing. To further minimize mixing of the electrolytes, a [[salt bridge]] can be used which consists of an electrolyte saturated gel in an inverted U-tube. As the negatively charged electrons flow in one direction around this circuit, the positively charged metal ions flow in the opposite direction in the electrolyte.
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| − | A [[galvanometer|voltmeter]] 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 Daniell 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 oxidized form by a vertical line, which represents the limit between the phases (oxidation changes). The double vertical lines represent the saline bridge on the cell. Finally, the oxidized form of the metal to be reduced at the cathode, is written, separated from its reduced form by the vertical line. The electrolyte concentration is given as it is an important variable in determining the cell potential.
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| − | ==Standard electrode potential==
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| − | {{Main|Standard electrode potential}}
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| − | To allow prediction of the cell potential, tabulations of [[standard electrode potential]] are available. Such tabulations are referenced to the standard hydrogen electrode (SHE). The [[standard hydrogen electrode]] undergoes the reaction
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| − | :<math>\mbox{2H}^{+}(aq) + \mbox{2e}^{-} \rightarrow \mbox{H}_{2}\,</math>
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| − | which is shown as reduction but, in fact, the SHE can act as either the anode or the cathode, depending on the relative oxidation/reduction potential of the other electrode/electrolyte combination. The term standard in SHE requires a supply of hydrogen gas bubbled through the electrolyte at a pressure of 1 atm and an acidic electrolyte with H+ activity equal to 1 (usually assumed to be [H+] = 1 mol/liter).
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| − | The SHE electrode can be connected to any other electrode by a salt bridge to form a cell. If the second electrode is also at standard conditions, then the measured cell potential is called the standard electrode potential for the electrode. The standard electrode potential for the SHE is zero, by definition. The polarity of the standard electrode potential provides information about the relative reduction potential of the electrode compared to the SHE. If the electrode has a positive potential with respect to the SHE, then that means it is a strongly reducing electrode which forces the SHE to be the anode (an example is Cu in aqueous CuSO4 with a standard electrode potential of 0.337 V). Conversely, if the measured potential is negative, the electrode is more oxidizing than the SHE (such as Zn in ZnSO4 where the standard electrode potential is -0.763 V).
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| − | Standard electrode potentials are usually tabulated as reduction potentials. However, the reactions are reversible and the role of a particular electrode in a cell depends on the relative oxidation/reduction potential of both electrodes. The oxidation potential for a particular electrode is just the negative of the reduction potential. A standard cell potential can be determined by looking up the standard electrode potentials for both electrodes (sometimes called half cell potentials). The one that is smaller will be the anode and will undergo oxidation. The cell potential is then calculated as the sum of the reduction potential for the cathode and the oxidation potential for the anode.
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| − | :<math>\mbox{E}^{o}_{cell}=\mbox{E}^{o}_{red}(cathode)-\mbox{E}^{o}_{red}(anode) = \mbox{E}^{o}_{red}(cathode)+\mbox{E}^{o}_{oxi}(anode) </math>
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| − | For example, the standard electrode potential for a copper electrode is:
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| − | :<math>\mbox{Cell diagram}\,</math>
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| − | :<math>\mbox{Pt}(s)|\mbox{H}_{2}(1 atm)|\mbox{H}^{+}(1 M)||\mbox{Cu}^{2+}(1 M)|\mbox{Cu}(s)\,</math>
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| − | :<math>\mbox{E}^{o}_{cell}=\mbox{E}^{o}_{red}(cathode)-\mbox{E}^{o}_{red}(anode)</math>
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| − | At standard temperature, pressure and concentration conditions, the cell's [[electromotive force|emf]] (measured by a [[multimeter]]) is 0.34 V. by definition, the electrode potential for the SHE is zero. Thus, the Cu is the cathode and the SHE is the anode giving
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| − | :<math>\mbox{E}_{cell}=\mbox{E}^{o}_{\mbox{Cu}^{2+}/\mbox{Cu}}-\mbox{E}^{o}_{\mbox{H}^{+}/\mbox{H}_{2}}</math>
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| − | Or,
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| − | :<math>\mbox{E}^{o}_{\mbox{Cu}^{2+}/\mbox{Cu}} = \mbox{0.34 V}</math>
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| − | Changes in the [[stoichiometric coefficient]]s of a balanced cell equation will not change <math>\mbox{E}^{0}_{red}\,</math> value because the standard electrode potential is an [[Intensive and extensive properties|intensive property]].
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| − | ==Spontaneity of Redox reaction==
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| − | {{main|Spontaneous process}}
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| − |
| |
| − | During operation of electrochemical cells, [[chemical energy]] is transformed into [[electrical energy]] and is expressed mathematically as the product of the cell's emf and the [[electrical charge]] transferred through the external circuit.
| |
| − | :<math>\mbox{Electrical energy}=\mbox{E}_{cell} \mbox{C}_{trans}\,</math>
| |
| − | where <math>\mbox{E}_{cell}\,</math> is the cell potential measured in volts (V) and <math>\mbox{C}_{trans}\,</math> is the cell current integrated over time and measured in coulumbs (C). <math>\mbox{C}_{trans}\,</math> can also be determined by multiplying the total number of electrons transferred (measured in moles) times Faraday's constant, F = 96,485 C/mole.
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| − |
| |
| − | The emf of the cell at zero current is the maximum possible emf. It is used to calculate the maximum possible electrical energy that could be obtained from a [[chemical reaction]]. This energy is referred to as [[electrical work]] and is expressed by the following equation:
| |
| − |
| |
| − | :<math>\mbox{W}_{max}=\mbox{W}_{electrical} = -\mbox{nFE}_{cell}\,</math>
| |
| − | where work is defined as positive into the system.
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| − |
| |
| − | Since the [[Thermodynamic free energy|free energy]] is the maximum amount of work that can be extracted from a system, one can write:
| |
| − | :<math>\Delta G=-\mbox{nFE}_{cell}\,</math>
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| − |
| |
| − | A positive cell potential gives a negative change in Gibbs free energy. This is consistent with the cell production of an electric current flowing from the cathode to the anode through the external circuit. If the current is driven in the opposite direction by imposing an external potential, then work is done on the cell to drive electrolysis.
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| − |
| |
| − | A [[spontaneous]] electrochemical reaction (change in Gibbs free energy less than zero) can be used to generate an
| |
| − | electric [[current (electricity)|current]], in [[electrochemical cell]]s. This is the basis of all 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, typically a combination of heat and electrical energy.
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| − |
| |
| − | 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.
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| − |
| |
| − | The relation between the [[equilibrium constant]], ''K'', and the Gibbs free energy for an electrochemical cell is expressed as follows:
| |
| − |
| |
| − | :<math>\Delta G^{o}=\mbox{-RT ln K}= \mbox{-nFE}^{o}_{cell}\,</math>
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| − |
| |
| − | Rearranging to express the relation between standard potential and equilibrium constant yields
| |
| − |
| |
| − | :<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>
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| − |
| |
| − | ==Cell emf dependency on changes in concentration==
| |
| − | ===Nernst Equation===
| |
| − | {{Main|Nernst Equation}}
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| − |
| |
| − | The standard potential of an electrochemical cell requires standard conditions for all of the reactants. When reactant concentrations differ from standard conditions, the cell potential will deviate from the standard potential. In the 20th century German [[chemist]] [[Walther Hermann Nernst]] proposed a mathematical model to determine the effect of reactant concentration on electrochemical cell potential.
| |
| − |
| |
| − | In the late 19th century [[Josiah Willard Gibbs]] had formulated a theory to predict whether a chemical reaction is spontaneous based on the free energy
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| − |
| |
| − | :<math>\Delta G=\Delta G^{o}+\mbox{RT ln Q}\,</math> ,
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| − |
| |
| − | Where:
| |
| − |
| |
| − | ''ΔG'' = change in [[Gibbs free energy]], ''T'' = absolute [[temperature]], ''R'' = [[gas constant]], ln = [[natural logarithm]], ''Q'' = [[reaction quotient]].
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| − |
| |
| − | Gibbs' key contribution was to formalize the understanding of the effect of reactant concentration on spontaneity.
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| − |
| |
| − | Based on Gibbs' work, Nernst extended the theory to include the contribution from electric potential on charged species. As shown in the previous section, the change in Gibbs free energy for an electrochemical cell can be related to the cell potential. Thus, Gibbs' theory becomes
| |
| − |
| |
| − | :<math>nF\Delta E = nF\Delta E^\circ - \mbox{RT ln Q} \,</math>
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| − |
| |
| − | Where:
| |
| − |
| |
| − | ''n'' = number of [[electrons]]/[[Mole (unit)|mole]] product, ''F'' = [[Faraday constant]] ([[coulomb]]s/mole), and ''ΔE'' = [[cell potential]].
| |
| − |
| |
| − | Finally, Nernst divided through by the amount of charge transferred to arrive at a new equation which now bears his name:
| |
| − | :<math>\Delta E=\Delta E^{o}- {\mbox{RT} \over \mbox{nF}} \mbox{ln Q}\,</math>
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| − |
| |
| − | Assuming standard conditions (<math>Temperature = 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>
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| − |
| |
| − | ===Concentration cells===
| |
| − | {{Main|Concentration cell}}
| |
| − | <!-- Deleted image removed: [[Image:Cell-membrane-electrochemical.PNG|thumb|Calculating [[membrane potential]] is a good example where concentration cells are used in biology to understand cellular [[metabolism]] such as the [[Na-K pump|Na<sup>+</sup>(red) K<sup>+</sup>(blue), or sodium-potassium pump]].]] —>
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| − | A concentration cell is an electrochemical cell where the two electrodes are the same material, the electrolytes on the two half-cells involve the same ions, but the electrolyte concentration differs between the two half-cells.
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| − |
| |
| − | For example an electrochemical cell, where two copper electrodes are submerged in two [[copper(II) sulfate]] solutions, whose concentrations are 0.05 [[Molar concentration|M]] and 2.0 M, connected through a salt bridge. This type of cell will generate a potential that can be predicted by the Nernst equation. Both electrodes undergo the same chemistry (although the reaction proceeds in reverse at the cathode)
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| − |
| |
| − | :<math>Cu^{2+}(aq)+2e^{-}\rightarrow \mbox{Cu}(s)</math>
| |
| − |
| |
| − | [[Le Chatelier's principle]] indicates that the reaction is more favourable to reduction as the concentration of <math>Cu^{2+}\,</math> ions increases. Reduction will take place in the cell's compartment where concentration is higher and oxidation will occur on the more dilute side.
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| − |
| |
| − | 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 the 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 the 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 the case mentioned, it is possible to calculate cell's potential:
| |
| − | :<math>E = 0- {0.0257 V \over 2} ln {0.05\over 2.0}= 0.0474{ } V\,</math>
| |
| − |
| |
| − | However, this value is only approximate, as reaction quotient is defined in terms of ion activities which can be approximated with the concentrations as calculated here.
| |
| − |
| |
| − | The Nernst equation plays an important role in understanding electrical effects in cells and organelles. Such effects include nerve [[synapses]] and [[cardiac cycle|cardiac beat]] as well as the resting potential of a somatic cell.
| |
| − |
| |
| − | ==Battery==
| |
| − | {{Main|Battery (electricity)}}
| |
| − |
| |
| − | A battery is an electrochemical cell (sometimes several in series) used for chemical energy storage. Batteries are optimized to produce a constant electric current for as long as possible. Although the cells discussed previously are useful for theoretical purposes and some laboratory experiments, the large internal resistance of the salt bridge make them inappropriate battery technologies. Various alternative battery technologies have been commercialized as discussed next.
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| − |
| |
| − | ===Dry cell===
| |
| − | {{Main|Dry cell}}
| |
| − | [[Image:Zincbattery.png|thumb|Zinc carbon battery diagram.]]
| |
| − | Dry cells do not have a [[fluid]] electrolyte. Instead, they use a moist electrolyte paste. [[Zinc-carbon battery|Leclanché's cell]] is a good example of this, where the 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]]. The cell's cathode is represented by a carbon bar inserted on the 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 around 1.5 [[Volt|V]].
| |
| − |
| |
| − | ===Mercury battery===
| |
| − | {{Main|Mercury battery}}
| |
| − | [[Image:Mercurybattery2.PNG|thumb|Cutaway view of a mercury battery.]]
| |
| − | The mercury battery has many applications in [[medicine]] and [[electronics]]. The battery consists of a [[steel]]—made container in the shape of a cylinder acting as the cathode, where an amalgamated anode of mercury and zinc is surrounded by a stronger alkaline electrolyte and a paste of [[zinc 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 in the electrolyte's composition when the cell works. Such batteries provide 1.35 V of [[direct current]].
| |
| − |
| |
| − | ===Lead-acid battery===
| |
| − | [[Image:Lead acid cell.jpg|thumb|A sealed lead-acid battery.]]
| |
| − | {{Main|Lead-acid battery}}
| |
| − |
| |
| − | The lead-acid battery used in [[automobiles]], consists of a series of six identical cells assembled in series. 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.
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| − |
| |
| − | 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>
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| − |
| |
| − | At standard conditions, each cell may produce a potential of 2 [[Volts|V]], hence overall voltage produced is 12 V. Differing from mercury and zinc-carbon batteries, lead-acid 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.
| |
| − |
| |
| − | ===Lithium rechargeable battery===
| |
| − | {{Main|Lithium battery}}
| |
| − | Instead of an [[aqueous]] electrolyte or a moist electrolyte paste, a solid state battery operates using a solid electrolyte. [[Lithium]] polymer batteries are an example of this; a graphite bar acts as the anode, a bar of lithium cobaltate acts as the [[cathode]], and a [[polymer]], swollen with a lithium salt, allows the passage of [[ions]] and serves as the electrolyte. In this cell, the carbon in the anode can reversibly form a lithium-carbon alloy. Upon discharging, lithium ions spontaneously leave the lithium cobaltate cathode and travel through the polymer and into the carbon anode forming the alloy. This flow of positive lithium ions is the electrical current that the battery provides. By charging the cell, the lithium dealloys and travels back into the cathode. The advantage of this kind of battery is that Lithium possess the highest negative value of standard reduction potential. It is also a [[light metal]] and therefore less mass is required to generate 1 [[faraday constant|mole of
| |
| − | electrons]]. Lithium ion battery technologies are widely used in portable electronic devices because they have high energy storage density and are rechargeable. These technologies show promise for future automotive applications, with new materials such as iron phosphates and lithium vanadates.
| |
| − |
| |
| − | ===Flow battery/ Redox flow battery===
| |
| − | {{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.
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| − |
| |
| − | 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 in [[power plants]] to supply electrical needs, however their conversion into electricity is an inefficient process. The most efficient electrical power plant may only convert about 40 percent of the original [[chemical energy]] into electricity when [[combustion|burned]] or processed.
| |
| − |
| |
| − | To enhance electrical production, scientists have developed fuel cells where [[combustion]] is replaced by electrochemical methods, similar to a battery but requiring continuous replenishment of the [[reactants]] consumed.
| |
| − |
| |
| − | The most popular is the oxygen-hydrogen fuel cell, where two inert electrodes ([[porous]] electrodes of [[nickel]] and [[nickel oxide]]) are placed in an electrolytic solution such as hot [[caustic potash]], in both compartments (anode and cathode) gaseous [[hydrogen]] and [[oxygen]] are bubbled into solution.
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| − |
| |
| − | Oxygen-hydrogen fuel cell reactions are shown bellow:
| |
| − | :<math>Anode: 2H_{2}(g)\rightarrow 4H^{+}+4e^{-}\,</math>
| |
| − | :<math>Cathode: O_{2}(g)+ 4e^{-} + 4 H^{+}\rightarrow 2H_{2}O(l)\,</math>
| |
| − | :<math>\mbox{Overall reaction:} 2H_{2}(g) + O_{2}(g)\rightarrow 2H_{2}O(l)\,</math>
| |
| − |
| |
| − | The overall reaction is identical to hydrogen [[combustion]]. Oxidation and reduction take place in the anode and [[cathode]] separately. This is similar to the electrode used in the cell for measuring standard reduction potential which has 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]], and [[rhodium]] are good electrocatalysts.
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| − |
| |
| − | ==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.
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| − |
| |
| − | ===Iron corrosion===
| |
| − | 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; [[hydronium|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. The concentration of water associated with iron oxide 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]].
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| − |
| |
| − | ===Corrosion of common metals===
| |
| − | [[Coinage metal]]s, such as copper and silver, slowly corrode through use.
| |
| − | A [[patina]] of green-blue [[copper carbonate]] forms on the surface of [[copper]] with exposure to the water and carbon dioxide in the air. [[Silver]] coins or [[cutlery]] that are exposed to high sulfur foods such as [[Egg (food)|egg]]s or the low levels of sulfur species in the air develop a layer of black [[Silver sulfide]].
| |
| − |
| |
| − | [[Gold]] and [[platinum]] are extremely difficult to oxidize under normal circumstances, and require exposure to a powerful chemical oxidizing agent such as [[aqua regia]].
| |
| − |
| |
| − | Some common metals oxidize extremely rapidly in air. [[Titanium]] and aluminium oxidize instantaneouly in contact with the oxygen in the air. These metals form an extremely thin layer of oxidized metal on the surface. This thin layer of oxide protects the underlying layers of the metal from the air preventing the entire metal from oxidizing. These metals are used in applications where corrosion resistance is important. [[Iron]], in contrast, has an oxide that forms in air and water, called [[rust]], that does not stop the further oxidation of the iron. Thus iron left exposed to air and water will continue to rust until all of the iron is oxidized.
| |
| − |
| |
| − | ===Prevention of corrosion===
| |
| − | Attempts to save a metal from becoming anodic are of two general types. Anodic regions dissolve and destroy the structural integrity of the metal.
| |
| − |
| |
| − | While it is almost impossible to prevent anode/[[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]]ed 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]]es exposing the metal will corrode with the region under the paint, adjacent to the scratch, to be anodic.
| |
| − |
| |
| − | Other prevention is called ''[[passivation]]'' where a metal is coated with another metal such as a [[tin can]]. Tin is a metal that rapidly corrodes to form a mono-molecular [[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 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, an ingot of buried or exposed magnesium (or zinc) is [[bury|buried]] beside the [[Pipe (material)|pipeline]] and is [[wire|connected electrically]] to the pipe above ground. The pipeline is forced to be a cathode and is protected from being oxidized and rusting. The magnesium anode is sacrificed. At intervals new [[ingot]]s are buried to replace those lost.
| |
| − |
| |
| − | ==Electrolysis==
| |
| − | {{Main|Electrolysis}}
| |
| − |
| |
| − | The spontaneous redox reactions of a conventional battery produce electricity through the different chemical potentials of the cathode and anode in the electrolyte. However, electrolysis requires an external source of [[electrical energy]] to induce a chemical reaction, and this process takes place in a compartment called an [[electrolytic cell]].
| |
| − |
| |
| − | ===Electrolysis of molten sodium chloride===
| |
| − |
| |
| − | When molten, the salt [[sodium chloride]] can be electrolyzed to yield metallic [[sodium]] and gaseous [[chlorine]]. Industrially, this process is carried out in a special cell called a Downs cell. The cell is connected to an electrical power supply, allowing [[electrons]] to migrate from the power supply to the electrolytic cell.
| |
| − |
| |
| − | Reactions that take place in the 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 large amounts of metallic sodium and gaseous chlorine, and is widely used on [[mineral dressing]] and [[metallurgy]] [[industry|industries]].
| |
| − |
| |
| − | The [[emf]] for this process is approximately -4 [[V]] indicating a (very) non-spontaneous process. In order for this reaction to occur the power supply should provide at least a potential of 4 V. However, larger voltages must be used for this reaction to occur at a high rate.
| |
| − |
| |
| − | ===Electrolysis of water===
| |
| − | {{Main|Electrolysis of water}}
| |
| − | Water can be converted to its component elemental gasses, H<sub>2</sub> and O<sub>2</sub> through the application of an external voltage. [[Water]] 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. The decomposition of water into hydrogen and oxygen can be performed in an electrolytic cell. In it, a pair of inert [[electrodes]] usually made of [[platinum]] immersed in water act as anode and cathode in the electrolytic process. The electrolysis starts with the application of an external voltage between the electrodes. This process will not occur except at extremely high voltages without an electrolyte such as [[sodium 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. While this reaction will work at any conductive electrode at a sufficiently large potential, platinum [[catalysis|catalyzes]] both hydrogen and oxygen formation, allowing for relatively mild voltages (~2V depending on the [[pH]]).
| |
| − |
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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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| − |
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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 with respect to its reduction and oxidation in both electrodes. Usually, water is electrolyzed, as mentioned above (in the 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]] into Na<sup>+</sup> and Cl<sup>-</sup> ions. The sodium ions are attracted to the cathode, where they are reduced to sodium metal. The chloride ions are attracted to the anode, where they are oxidized to chlorine gas.
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| − |
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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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| − |
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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<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.
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| − |
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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 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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| − |
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| − | Finally, reaction 3 is favorable because it describes the proliferation of [[hydroxide|OH<sup>-</sup>]] ions thus letting a probable reduction of [[hydronium|H<sup>+</sup>]] ions less favorable an option.
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| − |
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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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| − |
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| − | 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]].
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| − |
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| − | ===Quantitative electrolysis & Faraday's Laws===
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| − | {{Main|Faraday's law of electrolysis}}
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| − | Quantitative aspects of electrolysis were originally developed by [[Michael Faraday]] in 1834. Faraday is also credited with having coined the terms ''[[electrolyte]]'' and ''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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| − |
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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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| − |
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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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| − |
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| − | Below a simplified equation of Faraday's first law:
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| − | :<math>m \ = \ { 1 \over 96,485 \ \mathrm{(C \cdot mol^-1)} } \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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| − | :''M'' is the molar mass of the substance (in grams per [[mole (unit)|mole]]).
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| − |
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| − | ====Second law====
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| − | {{Main|Electroplating}}
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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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| − |
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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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| − |
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| − | ==Applications==
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| − |
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| − | There are various extremely important electrochemical processes in both nature and industry. They include the coating of objects with metals or metal oxides through electrodeposition, and the detection of alcohol in drunken drivers through the redox reaction of ethanol. The generation of chemical energy through [[photosynthesis]] is inherently an electrochemical process, as is the production of metals like aluminum and titanium from their ores. Certain diabetes blood sugar meters measure the amount of glucose in the blood through its redox potential.
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| − |
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| − | The nervous impulses in neurons are based on electric current generated by the movement of sodium and potassium ions into and out of cells. Some animals, such as eels, can generate a powerful voltage from certain cells that can disable much larger animals.
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| − |
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| − | == See also ==
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| − |
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| − | * [[Corrosion]]
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| − | * [[Electroplating]]
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| − | * [[Electricity]]
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| − | * [[Physical chemistry]]
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| − | * [[Redox]]
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| − |
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| − | ==References==
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| − |
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| − | * Bard, Allen J., György Inzelt, and F. Scholz, eds. 2008. ''Electrochemical Dictionary.'' Berlin: Springer. ISBN 9783540745976
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| − |
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| − | * Brown, Theodore E., H. Eugene LeMay, and Bruce E. Bursten. 2005. ''Chemistry: The Central Science,'' 10th ed. Upper Saddle River, NJ: Prentice Hall. ISBN 0131096869
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| − |
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| − | * Chang, Raymond. 2006. ''Chemistry''. 9th ed. New York: McGraw-Hill. ISBN 0073221031
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| − |
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| − | * Hill, John William, Ralph H. Petrucci, Terry McCreary, and Scott S. Perry. 2005. ''General Chemistry: An Integrated Approach.'' Upper Saddle River, NJ: Pearson Prentice Hall. ISBN 0131402838
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| − |
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| − | * Laidler, Keith James. 2001. ''The World of Physical Chemistry.'' Oxford: Oxford University Press. ISBN 0198559194
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| − |
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| − | * McMurry, John, and Robert C. Fay. 2004. ''Chemistry,'' 4th ed. Upper Saddle River, NJ: Prentice Hall. ISBN 0131402080.
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| − |
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| − | == External links ==
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| − |
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| − | * [http://scienceworld.wolfram.com/biography/Faraday.html Michael Faraday]. Eric Weisstein's World of Biography, Wolfram Research. Retrieved January 25, 2009.
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| − | * [http://library.thinkquest.org/19662/low/eng/electrolysis.html The Faraday Law of Electrolysis]. ThinkQuest. Retrieved January 25, 2009.
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| − | * [http://www.electrochem.org The Electrochemical Society]. Retrieved January 25, 2009.
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| − | * [http://electrochem.cwru.edu/estir/ Electrochemical Science and Technology Information Resource (ESTIR)]. Retrieved January 25, 2009.
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| − | * [http://www.ise-online.org International Society of Electrochemistry (ISE)]. Retrieved January 25, 2009.
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| − | * [http://electrochem.cwru.edu/ed/encycl/ Electrochemistry Encyclopedia]. Case Western Reserve University. Retrieved January 25, 2009.
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| − | * [http://electrochem.cwru.edu/ed/dict.htm Electrochemistry Dictionary]. Case Western Reserve University. Retrieved January 25, 2009.
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| − | * [http://www.funsci.com/fun3_en/electro/electro.htm Experiments in Electrochemistry]. Fun Science Gallery. Retrieved January 25, 2009.
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| − | {{BranchesofChemistry}}
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| − | {{Analytical chemistry}}
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| − | [[Category:Physical sciences]]
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| − | [[Category:Chemistry]]
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| − | [[Category:Physical chemistry]]
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| − | {{credit|242786656}}
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