electrochemistry: Definition and Much More from Answers.com

Wikipedia: electrochemistry

English chemists John Daniell (left) and Michael Faraday (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 conductor (the electrode composed of a metal or a semiconductor, including graphite) and an ionic conductor (the electrolyte).

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

History

Main article: History of electrochemistry

16th to 18th century developments

German physicist Otto von Guericke beside his electrical generator while conducting an experiment.

The 16th 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.

In 1663 the German physicist Otto von Guericke created the first electric generator, which produced static electricity by applying friction in the machine. The generator was made of a large sulfur ball cast inside a glass globe, mounted on a shaft. The ball was rotated by means of a crank and a static electric spark was produced when a pad was rubbed against the ball as it rotated. The globe could be removed and used as source for experiments with electricity.

By the mid—1700s the French chemist Charles François de Cisternay du Fay discovered two types of static electricity, and that like charges repel each other whilst unlike charges attract. Du Fay announced that electricity consisted of two fluids: "vitreous" (from the 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's one-fluid theory later in the century.

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.

Italian physicist Alessandro Volta showing his "battery" to French emperor Napoleon Bonaparte in early 1800s.

In the late 1700s the Italian physician and anatomist Luigi Galvani marked the birth of electrochemistry by establishing a bridge between chemical reactions and electricity on his essay "De Viribus Electricitatis in Motu Musculari Commentarius" (Latin for Commentary on the Effect of Electricity on Muscular Motion) in 1791 where he proposed a "nerveo-electrical substance" on biological life forms.

On his essay Galvani concluded that animal tissue contained a here-to-fore neglected innate, vital force, which he termed "animal electricity," which activated nerves and muscles spanned by metal probes. He believed that this new force was a form of electricity in addition to the "natural" form produced by lightning or by the electric eel and torpedo ray as well as the "artificial" form produced by friction (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

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 that the amount of metal deposited and the amount of oxygen produced during an electrolytic process depended on the distance between the electrodes. By 1801 Ritter observed thermoelectric currents and anticipated the discovery of thermoelectricity by Thomas Johann Seebeck.

By the 1810s William Hyde Wollaston made improvements to the galvanic pile. Sir Humphry Davy's work with electrolysis led to the conclusion that the production of electricity in simple electrolytic cells resulted from chemical action and that chemical combination occurred between substances of opposite charge. This work led directly to the isolation of sodium and potassium from their compounds and of the alkaline earth metals from theirs in 1808.

Hans Christian Ørsted's discovery of the magnetic effect of electrical currents in 1820 was immediately recognized as an epoch-making advance, although he left further work on electromagnetism to others. André-Marie Ampère quickly repeated Ørsted's experiment, and formulated them mathematically.

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 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 Daniell invented a primary cell in which hydrogen was eliminated in the generation of the electricity. Daniell had solved the problem of polarization. In his laboratory he had learned that alloying the amalgamated zinc of Sturgeon with mercury would produce a better voltage.

Swedish chemist Svante Arrhenius portrait circa 1880s.

William Grove produced the first fuel cell in 1839. In 1846, Wilhelm Weber developed the electrodynamometer. In 1866, Georges Leclanché patented a new cell which eventually became the forerunner to the world's first widely used battery, the zinc carbon cell.

Svante August Arrhenius published his thesis in 1884 on Recherches sur la conductibilité galvanique des électrolytes (Investigations on the galvanic conductivity of electrolytes). From his results the author concluded that electrolytes, when dissolved in water, become to varying degrees split or dissociated into electrically opposite positive and negative ions.

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

In 1894 Friedrich Ostwald concluded important studies of the electrical conductivity and electrolytic dissociation of organic acids.

German scientist Walther Nernst portrait in 1910s.

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 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–1970s quantum electrochemistry was developed by Revaz Dogonadze and his pupils.

Principles

Redox reactions

Main article: Redox reaction

Electrochemical processes are redox reactions where energy is produced by a spontaneous reaction which produces electricity, or where electrical current stimulates a chemical reaction. In a redox reaction, an atom's or ion's oxidation state (basically, its charge) changes as a result of an electron transfer.

Oxidation and Reduction

The elements involved in an electrochemical reaction are characterized by the number of electrons each has. The oxidation state of an ion is the number of electrons it has accepted or donated compared to its neutral state (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 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 article: Chemical equation

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

Acid medium

In acid medium H+ ions and water are added to half reactions to balance the overall reaction. For example, when Manganese reacts with Sodium bismuthate.

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

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

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

Finally the reaction is balanced by multiplying the number of electrons from the reduction half reaction to oxidation half reaction and vice versa and adding both half reactions, thus solving the equation.

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

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

Reaction balanced:

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

Basic medium

In basic medium OH^- ions and water are added to half reactions to balance the overall reaction. For example on reaction between Potassium permanganate and Sodium sulfite.

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

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

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

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

$\mbox{6e}^{-}+\mbox{4H}_{2}\mbox{O}+\mbox{2MnO}_{4}^{-}\rightarrow\mbox{2MnO}_{2}+\mbox{8OH}^{-}\,$

$\mbox{6OH}^{-}+\mbox{3SO}^{2-}_{3}\rightarrow\mbox{3SO}^{2-}_{4}+\mbox{3H}_{2}\mbox{O}+\mbox{6e}^{-}\,$

Equation balanced:

$\mbox{2KMnO}_{4}+\mbox{3Na}_{2}\mbox{SO}_3+\mbox{H}_2\mbox{O}\rightarrow\mbox{2MnO}_{2}+\mbox{3Na}_{2}\mbox{SO}_{4}+\mbox{2KOH}\,$

Neutral medium

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

$\mbox{Reaction unbalanced: }\mbox{C}_{3}\mbox{H}_{8}+\mbox{O}_{2}\rightarrow\mbox{CO}_{2}+\mbox{H}_{2}\mbox{O}\,$

$\mbox{Reduction: }\mbox{4H}^{+} + \mbox{O}_{2}+ \mbox{4e}^{-}\rightarrow\mbox{H}_{2}\mbox{O}+\mbox{H}_{2}\mbox{O}\,$

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

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

$\mbox{20H}^{+}+\mbox{5O}_{2}+\mbox{20e}^{-}\rightarrow\mbox{5H}_{2}\mbox{O}+\mbox{5H}_{2}\mbox{O}\,$

$\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\mbox{20e}^{-}+\mbox{20H}^{+}\,$

Equation balanced:

$\mbox{C}_{3}\mbox{H}_{8}+\mbox{5O}_{2}\rightarrow\mbox{3CO}_{2}+\mbox{4H}_{2}\mbox{O}\,$

Electrochemical cells

Main article: Electrochemical cell

A modified version of Daniells Cells, 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 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. 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 whose electrodes are Zinc and Copper submerged in Zinc sulfate and Copper sulfate, respectively, is known as a Daniells cell.

Half reactions for a Daniells cell are these:

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

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

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 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 called a salt bridge filled with an electrolytic solution is placed on the cell, thus allowing flow of ions, producing an electric current.

A voltmeter is capable of measuring the change of electrical potential between the anode and the cathode.

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

A cell diagram can be used to trace the path of the electrons in the electrochemical cell. For example, here is a cell diagram of a Daniells cell:

$\mbox{Zn}(s)|\mbox{Zn}^{2+}(1M)||\mbox{Cu}^{2+}(1M)|\mbox{Cu}(s)\,$

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.

Standard electrode potential

Main article: Standard electrode potential

Standard electrode potential is the value of the standard emf of a cell in which molecular hydrogen under standard pressure (10⁵ Pa) is oxidized to solvated protons at the left-hand electrode.

The cell potential depends on the difference between each half cell potential. Conventionally the potential associated with each electrode is chosen as the reduction takes place on the chosen electrode, hence standard electrode potential are tabulated on reduction potentials, thus tables are built on standard reduction potentials noted as $\mbox{E}^{0}_{red}\,$ .

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

$\mbox{E}^{o}_{cell}=\mbox{E}^{o}_{red}(cathode)-\mbox{E}^{o}_{red}(anode)$

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 $\mbox{E}^{0}_{red}=0\,$ . Cell's half reaction used for this procedure is hydrogen which in 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 (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, where there are plenty of H⁺ ions. Gaseous hydrogen enters through the tube and reacts over the platinum blade, thus allowing reduction and oxidation processes to occur.

The 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 the electric current generated on the cell represents the standard reduction potential for the element under measurement.

For example, on Copper standard reduction potential:

$\mbox{Cell diagram}\,$

$\mbox{Pt}(s)|\mbox{H}_{2}(1 atm)|\mbox{H}^{+}(1 M)||\mbox{Cu}^{2+}(1 M)|\mbox{Cu}(s)\,$

$\mbox{E}^{o}_{cell}=\mbox{E}^{o}_{red}(cathode)-\mbox{E}^{o}_{red}(anode)$

At standard temperature pressure conditions cell's emf (measured by a multimeter) is 0.34 V, conventionally SHE has a zero value, thus replacing on previous equation gives:

$\mbox{0.34V}_{cell}=\mbox{E}^{o}_{\mbox{Cu}^{2+}/\mbox{Cu}}-\mbox{E}^{o}_{\mbox{H}^{+}/\mbox{H}_{2}}$

$\mbox{0.34V}_{cell}=\mbox{E}^{o}_{\mbox{Cu}^{2+}/\mbox{Cu}}-0$

Electrochemical cell's emf value is used to predict whether redox reaction is a spontaneous process or not. A positive sign for overall cell's standard potential is considered to be spontaneous reaction, a negative sign would predict a spontaneous reaction on the opposite direction.

Changes over stoichiometric coefficients on balanced cell equation will not change $\mbox{E}^{0}_{red}\,$ value because standard electrode potential are intensive properties.

Spontaneity of Redox systems

Main article: Spontaneous process

On electrochemical cells, chemical energy transforms into electrical energy and is expressed mathematically as the product between the cell's emf by electrical charge in Coulombs.

$\mbox{Electrical energy}=(\mbox{volts})(\mbox{coulombs})\,$

$\mbox{Electrical energy}=\mbox{joules}\,$

The electrochemical cell's total charge is determined by multiplying the number of moles by Faraday's constant (F).

$\mbox{Total charge}=\mbox{n}\mbox{F}\,$

Faraday's constant is the electrical charge in 1 mole of electrons. It has been measured experimentally and is equivalent to 96 485.3 coulombs.

The cell's emf measured is the maximum voltage produced. It is used to calculate the maximum electrical energy obtained from a chemical reaction. This energy is referred to as electrical work and is expressed on the following equation:

$\mbox{W}_{max}=\mbox{W}_{electrical}\,$

$\mbox{W}_{max}=-\mbox{nFE}_{cell}\,$

Thus free energy is the amount of mechanical (or other) work that can be extracted from a system. Replacing this value on the previous equation with $\Delta G\,$ gives the relation between spontaneity and electrochemical cells.

$\Delta G=-\mbox{nFE}_{cell}\,$

The relation between Gibbs free energy and maximum electrical work may predict (at standard temperature and pressure conditions) whether the cell's redox system is a spontaneous process or not.

A spontaneous electrochemical reaction can be used to generate an electrical current, in electrochemical cells. This is the basis of all batteries and fuel cells. For example, gaseous oxygen (O₂) and hydrogen (H₂) can be combined in a fuel cell to form water and energy, typically a combination of heat and electrical energy.

Conversely, non-spontaneous electrochemical reactions can be driven forward by the application of a current at sufficient voltage. The 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:

$\Delta G^{o}=\mbox{-RT ln K}\,$

$\mbox{-nFE}^{o}_{cell}=\mbox{-RT ln K}\,$

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

$\mbox{E}^{o}_{cell}={\mbox{RT} \over \mbox{nF}} \mbox{ln K}\,$

Previous equation can use Briggsian logarithm as shown below:

$\mbox{E}^{o}_{cell}={0.0592 \mbox{V} \over \mbox{n}} \mbox{log K}\,$

Cell emf dependency on changes in concentration

Nernst Equation

Main article: Nernst Equation

Calculating a cell's potential is not always possible 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,

$\Delta G=\Delta G^{o}+\mbox{RT ln Q}\,$ ,

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 chemical activity.

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

$nF\Delta E = nF\Delta E^\circ - R T \ln Q \, \,$

Where:

n = number of electrons/mole product, F = Faraday constant (coulombs/mole), and ΔE = electrical potential of the reaction.

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

$\Delta E=\Delta E^{o}- {\mbox{RT} \over \mbox{nF}} \mbox{ln Q}\,$

Assuming standard conditions ( $Temperature = 298 K , 25 C\,$ ) and R = $8.3145 {J \over K mol}$ the equation above can be expressed on Base—10 logarithm as shown below:

$\Delta E=\Delta E^{o}- {\mbox{0.0592 V} \over \mbox{n}} \mbox{log Q}\,$

Concentration cells

Main article: Concentration cell

Calculating membrane potential is a good example where concentration cells are used in biology to understanding cellular metabolism such as the Na⁺(red) K⁺(blue), or sodium-potassium pump.

A concentration cell is an electrochemical cell whose electrodes are from the same material, but differing in ionic concentrations on both half-cells.

For example an electrochemical cell, where two copper electrodes are submerged on blue vitriol's solution, whose concentrations are 0.05 M and 2.0 M, connected through wire and saline bridge.

C u 2 + (a q) + 2 e - \toCu(s)

Le Chatelier's principle indicates reaction is favourable to reduction as the concentration of $Cu^{2+}\,$ ions increases. Reduction will take place in the cell's compartment where concentration is higher and oxidation will occur on the diluted side.

The following cell diagram describes the cell mentioned above:

$Cu(s)|Cu^{2+}(0.05 M)||Cu^{2+}(2.0 M)|Cu(s)\,$

Where the half cell reactions for oxidation and reduction are:

$Oxidation: Cu(s)\rightarrow \mbox{Cu}^{2+} (0.05 M) + 2e^{-}\,$

$Reduction: Cu^{2+} (2.0 M) +2e^{-} \rightarrow \mbox{Cu} (s)\,$

$Overall reaction: Cu^{2+} (2.0 M) \rightarrow \mbox{Cu}^{2+} (0.05 M)\,$

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

$E = E^{o}- {0.0257 V \over 2} ln {[Cu^{2+}]_{diluted}\over [Cu^{2+}]_{concentrated}}\,$

$E^{o}\,$ '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:

$E = 0- {0.0257 V \over 2} ln {0.05\over 2.0}\,$

$E = 0.0474 V\,$

However, this value is only approximate, as the potential difference is given from the ratio of activities of the ions, not the ratio of concentrations.

Concentration cells are often a significant matter of biological investigation; they are present on biological cells where membrane potential is responsible for nerve synapses and cardiac beat.

Battery

Main article: Battery (electricity)

A battery consists of one or more electrochemical cells, producing direct current at a constant voltage. The electrochemical principles are the same as electrochemical cells; however a battery does not need auxiliary components such as saline bridge on Daniell cells.

Dry cell

Main article: Dry cell

Zinc carbon battery diagram.

Dry cells do not have a fluid electrolyte. Instead, they use a moist electrolyte paste. 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 to have a pale and flabby consistency and avoiding flees. The cell's cathode is represented by a carbon bar inserted on the cell's electrolyte, usually placed in the middle.

Leclanché's simplified half reactions are shown below:

$Anode: Zn(s) \rightarrow Zn^{2+} (aq) + 2e^{-}\,$

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

$\mbox{Overall reaction:}\,$

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

The voltage obtained from the zinc-carbon battery is around 1.5 V.

Mercury battery

Main article: Mercury battery

Cutaway view of a Mercury battery diagram.

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: