Electrochemistry

Electrochemistry is a branch of electrolyte) and which involve electron transfer between the electrode and the electrolyte.

If a chemical reaction is caused by an external voltage, as in reduction reactions are separated in space, connected by an external electric circuit.

History

Main article: History of electrochemistry

16th to 18th century developments

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.

Charles-Augustin de Coulomb developed the law of electrostatic attraction in 1781 as an outgrowth of his attempt to investigate the law of electrical repulsions as stated by Joseph Priestley in England. In the late 1700s the Italian physician and anatomist Luigi Galvani marked the birth of electrochemistry by establishing a bridge between chemical reactions and electricity on his essay "De Viribus Electricitatis in Motu Musculari Commentarius" (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

In 1800, the English chemists William Nicholson and Johann Wilhelm Ritter succeeded in decomposing water into 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 alkaline earth metals from theirs in 1808.

Hans Christian Ørsted's discovery of the magnetic effect of electrical currents in 1820 was immediately recognized as an epoch-making advance, although he left further work on electromagnetism to others. André-Marie Ampère quickly repeated Ørsted's experiment, and formulated them mathematically. In 1821, Estonian-German physicist Thomas Johann Seebeck demonstrated the electrical potential in the juncture points of two dissimilar metals when there is a heat difference between the joints.

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

In 1832 mercury would produce a better voltage. zinc carbon cell.

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 organic acids. Nernst Equation, which related the voltage of a cell to its properties.

In 1898 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, Thomas Martin Lowry published essentially the same theory about how acids and bases behave, using an electrochemical basis.

electrophoretic apparatus in 1937 and some years later he was awarded 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 Revaz Dogonadze and his pupils.

Principles

Redox reactions

Main article: Redox reaction

Electrochemical processes involve 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 electron transfer.

Oxidation and Reduction

The redox reaction.

For example when ionic bond.

The loss of electrons from an atom or molecule is called 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.

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 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: electronegativity) than oxygen.

For reactions involving oxygen, the gain of oxygen implies the oxidation of the atom or molecule to which the oxygen is added (the oxygen is reduced). This follows since oxygen is highly electronegative (this can be extended to include other highly electronegative atoms such a fluorine). In compounds involving non-metals, the loss of hydrogen implies oxidation of the atom or molecule from which it is lost (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.

Electrochemical cells

Main article: Electrochemical cell

An electrochemical cell is a device capable of producing electric current from energy released by a 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.

Electrochemical cells have two electrodes (the anode and the cathode). The cathode is the electrode where the reduction takes place. Electrodes come in various forms including metal, gas and carbon.

The Galvanic cell uses two different metal electrodes, each in an electrolyte where the cations are the oxidized form of the electrode metal. The tendency of the electrode metals to oxidize or reduce, in a particular electrolyte, is defined in terms of the electrochemical potential which depends on the temperature, pressure, and the componsition and concentration of the electrolyte. In a particular cell, one electrode will undergo oxidation (the anode) and the other will undergo reduction (the cathode), depending on the sum of the electrochemical potential differences at both electrodes. Thus, a cell will produce an electrical potential difference between the two electrodes whose magnitude and sign depends on the chemistry occurring at both electrodes.

A Galvanic cell whose copper sulfate, respectively, is known as a Daniells cell.

Half reactions for a Daniells cell are these:

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

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

To provide a complete electric circuit, there must be a conduction path between the anode and cathode electrolytes. The simplest ionic conduction path is to provide a liquid junction. To avoid convective mixing between the two electrolytes, the liquid junction can be provided through a porous plug that allows ion flow while restricting convective flow. Such porous plug connections allow the anions and cations from the two electrolytes to mix, which can cause contamination. To minimize such issues, a salt bridge can be used which consists of an electrolyte saturated gel in an inverted U-tube. Contamination issues can be minimized by careful choices of the electrolytes in the anode, cathode and salt bridge.

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. The electrolyte concentration is given as it is an important variable in determining the cell potential.

Standard electrode potential

Main article: Standard electrode potential

To allow prediction of the cell potential, tabulations of standard hydrogen electrode undergoes the reaction

$\mbox{2H}^{+}(aq) + \mbox{2e}^{-} \rightarrow \mbox{H}_{2}\,$

which is shown as reduction but, in fact, the SHE can act as either the anode or the cathode, depending on the relative oxidation/reduction potential of the other electrode/electrolyte combination. The term standard in SHE requires a supply of hydrogen gas bubbled through the electrolyte at a pressure of 1 atm and an acidic electrolyte with H+ activity equal to 1 (usually assumed to be [H+] = 1 mol/liter).

The SHE electrode can be connected to any other electrode by a salt bridge to form a cell. If the second electrode is also at standard conditions, then the measured cell potential is called the standard electrode potential for the electrode. The standard electrode potential for the SHE is zero, by definition. The polarity of the standard electrode potential provides information about the relative reduction potential of the electrode compared to the SHE. If the electrode has a positive potential with respect to the SHE, then that means it is a strongly reducing electrode which forces the SHE to be the anode (an example is Cu in aqueous CuSO4 with a standard electrode potential of 0.337 V). Conversely, if the measured potential is negative, the electrode is more oxidizing than the SHE (such as Zn in ZnSO4 where the standard electrode potential is -0.763 V).

Standard electrode potentials are usually tabulated as reduction potentials. However, the reactions are reversible and the role of a particular electrode in a cell depends on the relative oxidation/reduction potential of both electrodes. The oxidation potential for a particular electrode is just the negative of the reduction potential. A standard cell potential can be determined by looking up the standard electrode potentials for both electrodes (sometimes called half cell potentials). The one that is smaller will be the anode and will undergo oxidation. The cell potential is then calculated as the sum of the reduction potential for the cathode and the oxidation potential for the anode.

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

For example, the standard electrode potential for a copper electrode is:

$\mbox{Cell diagram}\,$

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

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

At standard temperature, pressure and concentration conditions, the cell's emf (measured by a multimeter) is 0.34 V. by definition, the electrode potential for the SHE is zero. Thus, the Cu is the cathode and the SHE is the anode giving

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

Or,

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

Changes in the stoichiometric coefficients of a balanced cell equation will not change $\mbox{E}^{0}_{red}\,$ value because the standard electrode potential is an intensive property.

Spontaneity of Redox reaction

Main article: Spontaneous process

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.

$\mbox{Electrical energy}=\mbox{E}_{cell} \mbox{C}_{trans}\,$

where $\mbox{E}_{cell}\,$ is the cell potential measured in volts (V) and $\mbox{C}_{trans}\,$ is the cell current integrated over time and measured in coulumbs (C). $\mbox{C}_{trans}\,$ can also be determined by multiplying the total number of electrons transferred (measured in moles) times Faraday's constant, F = 96,485 C/mole.

The emf of the cell at zero current is the maximum possible emf. It is used to calculate the maximum possible electrical energy that could be obtained from a chemical reaction. This energy is referred as electrical work and is expressed by the following equation:

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

where work is defined as positive into the system.

Since the free energy is the maximum amount of work that can be extracted from a system, one can write:

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

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.

A spontaneous electrochemical reaction (change in Gibbs free energy less than zero) can be used to generate an electric current, in 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 the equilibrium constant and the Gibbs free energy for an electrochemical cell is expressed as follows:

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

Rearranging to express the relation between standard potential and equilibrium constant yields

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

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

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 1900s German Walther Hermann Nernst proposed a mathematical model to determine the effect of reactant concentration on electrochemical cell potential.

In the late 1800s Willard Gibbs formulated a theory to predict whether a chemical reaction is spontaneous based on the free energy

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

Where:

ΔG = change in reaction quotient.

Gibbs key contribution was to formalize the understanding of the effect of reactant concentration on spontaneity.

Based on Gibbs's 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

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

Where:

n = number of 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:

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

Assuming standard conditions ( $Temperature = 25 C\,$ ) 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

A concentration cell is an electrochemical cell where the two electrodes are the same material, the electrolytes on the two half-cells involve the same ions, but the electrolyte concentration differs between the two half-cells.

For example an electrochemical cell, where two copper electrodes are submerged in blue vitriol's solution, whose concentrations are 0.05 M and 2.0 M, connected through a salt bridge. This type of cell will generate a potential that can be predicted by the Nernst equation. Both electrodes undergo the same basic chemistry (although the reaction proceeds in reverse at the cathode)

$Cu^{2+}(aq)+2e^{-}\rightarrow \mbox{Cu}(s)$

Le Chatelier's principle indicates that the reaction is more 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 more dilute 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:

$E = 0- {0.0257 V \over 2} ln {0.05\over 2.0}= 0.0474{ } V\,$

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 beat as well as the resting potential of a somatic cell.

Battery

Main article: 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, including a salt bridge, are useful for theoretical purposes and some laboratory purposes, the large internal resistance of the salt bridge make them inappropriate battery technologies. Various alternative battery technologies have been commercialized as discussed next.

Dry cell

Main article: Dry cell

Dry cells do not have a starch. The cell's cathode is represented by a carbon bar inserted on the cell's electrolyte, usually placed in the middle.

Leclanché's simplified half reactions are shown below:

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

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

$\mbox{Overall reaction:}\,$

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

The voltage obtained from the V.

Mercury battery

Main article: Mercury battery

The mercury battery has many applications in medicine and electronics. The battery consists of a Zinc oxide and Mercury(II) oxide .

Mercury battery half reactions are shown below:

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

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

$\mbox{Overall reaction:}\,$

$Zn(Hg) + HgO(s) \rightarrow ZnO(s) + Hg(l)\,$

There are no changes in the electrolyte's composition when the cell works. Such Mercurium batteries provide 1.35 V of direct current.

Lead-acid battery

Main article: 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 sulfuric acid acting as the electrolyte.

Lead-acid battery half cell reactions are shown below:

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

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

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

At standard conditions, each cell may produce a potential of 2 electrolysis of the products in the overall reaction (discharge), thus recovering initial components which made the battery work.

Solid state Lithium battery

Main article: Lithium battery

Instead of an aqueous electrolyte or a moist electrolyte paste, a solid state battery operates using a solid electrolyte. Solid state light metal and therefore less mass is required to generate 1 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.

Flow battery/ Redox flow battery

Main article: Flow battery

Most batteries have all of the electrolyte and electrodes within a single housing. A flow battery is unusual in that the majority of the electrolyte, including dissolved reactive species, is stored in separate tanks. The electrolytes are pumped through a reactor, which houses the electrodes, when the battery is charged or discharged.

These types of batteries are typically used for large-scale energy storage (kWh - multi MWh). Of the several different types that have been developed, some are of current commercial interest, including the zinc bromine battery.

Fuel cells

Main article: Fuel cell

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% of the original chemical energy into electricity when 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 oxygen are bubbled into solution.

Oxygen-hydrogen fuel cell reactions are shown bellow:

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

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

$\mbox{Overall reaction:} 2H_{2}(g) + O_{2}(g)\rightarrow 2H_{2}O(l)\,$

The overall reaction is identical to rhodium are good electrocatalysts.

Corrosion

Main article: Corrosion

Corrosion is the term applied to brass. The cost of replacing metals lost to corrosion is in the multi-billions of dollars per year.

Iron corrosion

For iron rust to occur the metal has to be in contact with oxygen and water, although chemical reactions for this process are relatively complex and not all of them are completely understood, it is believed the causes are the following:

Electron transferring (Reduction-Oxidation)
1. 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.
  1. $Fe(s)\rightarrow Fe^{2+}(aq) + 2e^{-}\,$
2. water on the cathode, which is placed in another region of the metal.
  1. $O_{2}(g) + 4H^{+}(aq) + 4e^{-} \rightarrow 2H_{2}O(l)\,$
3. Global reaction for the process:
  $2Fe(s) + O_{2}(g) + 4H^{+}(aq) \rightarrow 2Fe^{2+}(aq) + 2H_{2}O(l)\,$
4. Standard emf for iron rusting:
  1. $E^{o}=E^{o}_{cathode}-E^{o}_{anode}\,$
    $E^{o}=1.23V-(-0.44V)=1.67V\,$

Iron corrosion takes place on acid medium; carbonic acid. Fe²⁺ ions oxides, following this equation:

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

hydrated is known as rust. Water associated with iron oxide it varies, thus chemical representation is presented as $Fe_{2}O_{3}.xH_{2}O\,$ . The electric circuit works as passage of electrons and ions occurs, thus if an electrolyte is present it will facilitate salt water.

Corrosion of coinage metals

Coinage metals, such as copper and silver, can also slowly corrode. At standard temperature and pressure, a Silver cutlery that is in contact with food can develop a layer of Silver sulfide.

Prevention of Corrosion

Attempts to save a metal from becoming anodic are of two general types. Anodic regions dissolve and destroy the structural integrity of the metal.

While it is almost impossible to prevent electrolyte is not possible and corrosion will not occur.

Coating

Metals are coated on its surface with anodic.

Other prevention is called iron). However, if the tin coating is scratched the iron becomes anodic and the can corrodes rapidly.

Sacrificial anodes

A method commonly used to protect a structural metal is to attach a metal which is more anodic than the metal to be protected. This forces the structural metal to be cathodic, thus spared corrosion. It is called "sacrificial" because the anode dissolves and has to be replaced periodically.
magnesium, would work very well but zinc is the least expensive useful metal.
To protect pipelines, buried or exposed an ingot of magnesium (or zinc) is buried beside the pipeline and connected electrically to the pipe above ground. The pipeline is forced to be a cathode and is protected. The magnesium anode is sacrificed. At intervals new ingots are buried to replace those lost.

Electrolysis

Main article: Electrolysis

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

Electrolysis of molten sodium chloride

When molten, electrons migration from the battery to the electrolytic cell.
Reactions that take place at Down's cell are the following:

$\mbox{Anode (oxidation): }2Cl^{-} \rightarrow Cl_{2}(g) + 2e^{-}\,$
$\mbox{Cathode (reduction): }2Na^{+}(l) + 2e^{-} \rightarrow 2Na(l)\,$
$\mbox{Overall reaction: }2Na^{+} + 2Cl^{-}(l) \rightarrow 2Na(l) + Cl_{2}(g)\,$

This process can yield industrial amounts of metallic sodium and gaseous chlorine, and is widely used on mineral dressing and metallurgy industries.
Standard emf for this process is approximately -4 V indicating a non-spontaneous process. In order this reaction to occur the battery should provide at least a potential of 4V. However, on mineral refining industry, higher voltages are used, due to low efficiency of the process.

Electrolysis of water

Main article: Electrolysis of water

Water at standard temperature and pressure conditions doesn't decompose into Gibbs free energy for the process at standard conditions is about 474.4 kJ
However, special sulfuric acid (most used 0.1 M).
Bubbles from the gases will be seen near both electrodes. The following half reactions describe the process mentioned above:

$\mbox{Anode (oxidation): }2H_{2}O(l) \rightarrow O_{2}(g) + 4H^{+}(aq) + 4e^{-}\,$
$\mbox{Cathode (reduction): }2H_{2}O(g) + 2e^{-} \rightarrow H_{2}(g) + 2OH^{-}(aq)\,$
$\mbox{Overall reaction: }2H_{2}O(l) \rightarrow 2H_{2}(g) + O_{2}(g)\,$

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

Electrolysis of aqueous solutions

Electrolysis in an aqueous is a similar process as mentioned in electrolysis of water. However, it is considered to be a complex process because the contents in solution have to be analyzed in half reactions, whether reduced or oxidized.

Electrolysis of a solution of Sodium chloride

The presence of water in a solution of chloride ion.
The following half reactions describes the process mentioned:

$\mbox{1. Cathode: }Na^{+}(aq)+ 1e^{-} \rightarrow Na(s) \qquad E^{o}_{red}=-2.71 V\,$
$\mbox{2. Anode: }2Cl^{-}(aq) \rightarrow Cl_{2}(g) + 2e^{-} \qquad E^{o}_{red}= +1.36 V\,$
$\mbox{3. Cathode: }2H_{2}O(l) + 2e^{+} \rightarrow H_{2}(g) + 2OH^{-}(aq)\qquad E^{o}_{red}=-0.83 V\,$
$\mbox{4. Anode: } 2H_{2}O(l) \rightarrow O_{2}(g) + 4H^{+}(aq) + 4e^{-}\qquad E^{o}_{red}=+1.23V\,$

Reaction 1 is discarded as it has the most negative value on standard reduction potential thus making it less thermodynamically favorable in the process.
When comparing the reduction potentials in reactions 2 & 4, the reduction of chloride ion is favored. Thus, if the Cl^- ion is favored for oxidation producing gaseous oxygen, however experiments shown gaseous chlorine is produced and not oxygen.
Although the initial analysis is correct, there is another effect that can happen, known as the overvoltage effect. Additional voltage is sometimes required, beyond the voltage predicted by the $E^{o}_{cell}\,$ . This may be due to kinetic terms. In other words, although the voltage applied is thermodynamically sufficient to drive electrolysis, the rate is so slow that to make the process proceed in a reasonable time frame, the voltage of the external source has to be increased (hence, overvoltage).
Finally, reaction 3 is favorable because it describes the proliferation of H⁺ ions less favorable an option.
The overall reaction for the process according to the analysis would be the following:

$\mbox{Anode (Oxidation): } 2Cl^{-}(aq)\rightarrow Cl_{2}(g) + 2e^{-}\,$
$\mbox{Cathode (Reduction): } 2H_{2}O(l) + 2e{-}\rightarrow H_{2}(g) + 2OH^{-}(aq)\,$
$\mbox{Overall reaction: } 2H_{2}O + 2Cl^{-}(aq) \rightarrow H_{2}(g) + Cl_{2}(g) + 2OH^{-}(aq)\,$

As the overall reaction indicates, the sodium hydroxide.

Quantitative electrolysis & Faraday Laws

Main article: Faraday's law of electrolysis

Quantitative aspects of electrolysis were originally developed by electrolyte, electrolysis, among many others while he studied quantitative analysis of electrochemical reactions. Also he was an advocate of the law of conservation of energy.

First law

Faraday concluded after several experiments on electrical current in non-spontaneous process, the mass of the products yielded on the electrodes was proportional to the value of current supplied to the cell, the length of time the current existed, and the molar mass of the substance analyzed.

In other words, the amount of a substance deposited on each electrode of an electrolytic cell is directly proportional to the quantity of electricity passed through the cell.

Below a simplified equation of Faraday's first law:

$m \ = \ { 1 \over 96,485 \ \mathrm{(C \cdot mol^-1)} } \cdot { Q M \over n }$

Where,

m is the mass of the substance produced at the electrode (in grams),

Q is the total electric charge that passed through the solution (in coulombs),

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

M is the molar mass of the substance (in grams per mole).

Second law

Main article: Electroplating

Faraday devised the laws of chemical electrodeposition of metals from solutions in 1857. He formulated the second law of electrolysis stating "the amounts of bodies which are equivalent to each other in their ordinary chemical action have equal quantities of electricity naturally associated with them." In other terms, the quantities of different elements deposited by a given amount of electricity are in the ratio of their chemical equivalent weights.

An important aspect of the second law of electrolysis is corrosion.

Applications

There are various electrochemical processes in both nature and industry, like the fabrication of devices with metals or aleations covered with thin deposits of other metal to prevent these artifacts from having corrosion, the reproduction of objects in copper through electrolysis, and the detection of ingested liquor in drunken drivers through the redox reaction.

The nervous impulses in neurons are based on electric charges with sodium ions and the expulsion of potassium ions from the cell; also, the body of an electric eel has cells that, individually, sponsor small electric charges, but together produce a 700-volt potential, enough to kill an human or a horse.

References

Electrochemistry. General Chemistry II by Dr. Michael Blaber. Retrieved on January 30, 2006.
Electrochemistry. Corrosion. Retrieved on January 28, 2006.
Michael Faraday. Biography. Retrieved on January 30, 2006.
The Faraday law of electrochemistry. Faraday laws of electrochemistry. Retrieved on January 30, 2006.
Chang, Raymond (2002). "Electrochemistry", Chemistry, 7th Edition, Mc Graw Hill. ISBN 0-07-365601-1.
L. Brown, Theodore; H. Eugene LeMay, Jr., Bruce E. Bursten, Julia R. Burdge (2003). "Electrochemistry", Chemistry, 9th Edition, US: Pearson Education. ISBN 0-13-066997-0.
William Hill, John; Ralph H. Petrucci, Terry McCreary, Scott S. Perry (March 2004). "Electrochemistry", General Chemistry: An Integrated Approach, 7th Edition, Pearson Education, 1200. ISBN 0-13-140283-8.
McMurry, John; Robert C. Fay (March 2004). "Electrochemistry", Chemistry, 3rd Edition, Prentice Hall. ISBN 0-13-056765-5.
Laidler, Keith; John H. Meiser, Bryan C. Sanctuary (May 2002). "Electrochemistry", in Keith Laidler: Physical Chemistry, 4th Edition, Boston, U.S.: Houghton Mifflin Company College Division. ISBN 061815292X.

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Categories: Physical chemistry

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