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Faraday-Daniell

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

Guericke-electricaldevice

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

The 16th century marked the beginning of the electrical understanding. During the 1550s the English scientist William Gilbert spent 17 years experimenting with magnetism and, to a lesser extent, electricity. For his work on magnets, Gilbert became known as the "Father of Magnetism." He discovered various methods for producing and strengthening magnets.

In 1663 the 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.
Galvani-frog-legs

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.

Volta-and-napoleon

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

Humphrydavy

Sir Humphry Davy's portrait in 1800s.

In 1800, the English chemists William Nicholson (chemist) and Johann Ritter succeeded in decomposing water into hydrogen and oxygen by electrolysis. Soon thereafter Johann Ritter discovered the process of electroplating. He also observed the amount of metal deposited and the amount of oxygen produced during an electrolytic process that depended on the distance between the electrodes. By 1801 Ritter observed 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.

ChemicalHistoryofaCandle

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.

Arrhenius2

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.

Walther Nernst 2

German scientist Walther Nernst portrait in 1910s.

Hermann Nernst developed the theory of the electromotive force of the voltaic cell in 1888. In 1889, he showed how the characteristics of the current produced could be used to calculate the free energy change in the chemical reaction producing the current. He constructed an equation, known as Nernst Equation, which related the voltage of a cell to its properties.

In 1898 Fritz Haber showed that definite reduction products can result from electrolytic processes if the potential at the cathode is kept constant. In 1898 he explained the reduction of nitrobenzene in stages at the cathode and this became the model for other similar reduction processes.


The 20th century and recent developments

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

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

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

Arne Tiselius developed the first sophisticated electrophoretic apparatus in 1937 and some years later he was awarded to the 1948 Nobel Prize for his work in protein electrophoresis.

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

By the 1960s1970s quantum electrochemistry was developed by Revaz Dogonadze and his pupils.

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, sodium chloride can be electrolysed to yield metallic sodium and gaseous chlorine. Industrially this process takes place in a special cell named Down's cell. The cell is connected to a battery, allowing electrons 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 on the process.

Electrolysis of water

Hoffman voltameter

Diagram of a Hofmann voltameter, showing electrolysis of water.

Main article: Electrolysis of water

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

However, special laboratory glassware has been designed for this purpose- the Hofmann voltameter. In it, a pair of inert electrodes usually made of platinum act as anode and cathode in the electrolytic process. After the water (if pure) has been placed in the apparatus, nothing happens, hence there are not enough ions to let the passage of electrons occur. To start the electrolysis an electrolyte should be placed in, usually sodium chloride or sulfuric acid (most used 0.1 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 sodium chloride must be examined in respect to its reduction and oxidation in both electrodes. Usually, water is electrolysed as mentioned in electrolysis of water yielding gaseous oxygen in the anode and gaseous hydrogen in the cathode. On the other hand, sodium chloride in water dissociates in Na+ and Cl- ions, anion will be attracted to the cathode, thus reducing the sodium ion. The cation will then be attracted to the anode oxidizing chloride ion.

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 reduction, then the water reaction is favored for oxidation producing gaseous oxygen, however experiments shown gaseous chlorine is produced and not oxygen.

Although the initial analysis is correct, there is another effect that can happen, known as the overvoltage effect. Additional voltage is sometimes required, beyond the voltage predicted by the E^{o}_{cell}\,. This may be due to kinetic rather than 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).

Finally, reaction 3 is favorable because it describes the proliferation of OH- ions thus letting a probable reduction 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 concentration of chloride ions is reduced in comparison to OH- ions (whose concentration increases). The reaction also shows the production of gaseous hydrogen, chlorine and aqueous sodium hydroxide.

Quantitative electrolysis & Faraday Laws

Main article: Faraday's law of electrolysis

Quantitative aspects of electrolysis were originally developed by Michael Faraday in 1834. Faraday is also credited to have coined the terms electrolyte, electrolysis, among many others while he studied quantitative analysis of electrochemical reactions. Also he was an advocate of the law of conservation of energy.

File:Nickel-electroplating.PNG

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 F \ } \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),
F is the Faraday's constant = 96,485 coulombs per mole,
M is the molar mass of the substance (in grams per mole).

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

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 electroplating which together with the first law of electrolysis, has a significant number of applications in the industry, as when used to protect metals to avoid corrosion.

Principles

Redox reactions

Main article: Redox reaction

Electrochemical process are redox reactions where energy is produced by a spontaneous reaction which produces electricity, otherwise electrical current stimulates a chemical reaction. In a redox reaction, an atom's oxidation state changes as a result of an electron transfer.

Oxidation and Reduction

The 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 atoms 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}\rightarrow\mbox{H}_{2}\mbox{O}+\mbox{H}_{2}\mbox{O}+ \mbox{4e}^{-}\,
\mbox{Oxidation: }\mbox{20e}^{-}+\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\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}\rightarrow\mbox{5H}_{2}\mbox{O}+\mbox{5H}_{2}\mbox{O}+\mbox{20e}^{-}\,
\mbox{20e}^{-}+\mbox{6H}_{2}\mbox{O}+\mbox{C}_{3}\mbox{H}_{8}\rightarrow\mbox{3CO}_{2}+\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
Galvanic 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 which electrodes are Zinc and Copper submerged on Zinc sulfate and Copper sulfate respectively is known as 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)\,
BASi epsilon C3 cell stand

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 filled with an electrolytic solution is placed on the cell, thus allowing flow of electrons, producing D.C. electric current.

A voltameter 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 oxidised form by a vertical line, which represents the limit between the phases (oxidation changes). The double vertical lines represent the saline bridge on the cell. Finally, the oxidized form of the metal to be reduced at the cathode, is written, separated from its reduced form by the vertical line.

Standard electrode potential

Main article: Standard electrode potential
File:Reduction-potentials2.PNG

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, plenty of H+ ions, gaseous hydrogen enter through the tube and react over the platinum blade thus allowing reduction and oxidation processes to occur.

SHE operates exactly as the same way as conventional electrodes on Daniells cell's work; in order to measure the standard reduction potential, SHE replaces one of the electrodes in the electrochemical cell acting as cathode or anode, thus electric current generated on the cell represents the standard reduction potential for the element which is measured.

For example on Copper standard reduction potential:

\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 cell's emf by electrical charge in Coulombs.

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

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.

Cell's emf measured is the maximum voltage produced, this value is used to calculate the maximum electrical energy which is obtained from a chemical reaction, this energy is referred to as electrical work and is expressed on the following equation,

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

Conversely, non-spontaneous electrochemical reactions can be driven forward by the application of a current at sufficient voltage. The electrolysis of water into gaseous oxygen and hydrogen is a typical example.

The relation between equilibrium constant and spontaneity based on Gibbs free energy terms on electrochemical cells is expressed as follows:

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

Solving both equations express 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 cell's potential is not always plausible at standard temperature and pressure conditions. However in 1900s German chemist Walther Hermann Nernst proposed a mathematical model to determine electrochemical cell potential where standard conditions cannot be reached.

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

\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 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 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
Cell-membrane-electrochemical

Calculating membrane potential is good example where concentration cells are used in biology to understanding cell's metabolism such as Na+(red) K+(blue) pump.

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

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

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

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

The following cell diagram describes the cell mentioned above:

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

Where both 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 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 case mentioned 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, because the potential difference is given from the ratio of activities of the ions, not the ratio of concentrations.

Concentration cell's are often a significant biologist's matter of investigation hence they are present on biological cells where membrane potential is responsible of nerve synapses and cardiac beat.

Battery

Main article: Battery (electricity)

A battery is an electrochemical cell or a group of them, where if combined together, may produce direct current at a constant voltage. Electrochemical principles which made batteries work are the same as on electrochemical cells, however a battery doesn't need auxiliary components such as saline bridge on Daniell cells.

Dry cell

Main article: Dry cell
Zincbattery

Zinc carbon battery diagram.

Dry cells don't have a fluid electrolyte instead they use a moist electrolyte paste. Leclanché's cell is a good example of this, where cell's anode is a zinc container surrounded by a thin layer of manganese dioxide and a moist electrolyte paste of ammonium chloride and zinc chloride mixed with starch to have a pale and flabby consistency and avoiding flees. Cell's cathode is represented by a carbon bar inserted on cell's electrolyte, usually placed in the middle.

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 1.5 V approximately.

Mercury battery

Main article: Mercury battery
Mercurybattery2

Cutaway view of a Mercury battery diagram.

Mercury battery has many applications on medicine and electronics. The battery consists of a steel—made container with the shape of a cylinder acting as the cathode, where an 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:

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 on the electrolyte's composition when cell works. Mercurium battery provides 1.35 V of direct current.

Lead-acid battery

Lead acid cell

A sealed Lead acid battery.

Main article: Lead-acid battery

The Lead-acid battery used on automobiles, consists on a series of six identical cells in line assembled, each cell has a lead anode and a cathode made from lead dioxide packed in a metal plaque. Cathode and anode are submerged in a solution of sulfuric acid acting as the electrolyte.

Lead-acid battery half cell reactions are shown below:

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 direct current of 2 V, hence overall voltage produced is 12 V. Lead-acid batteries, differing from Mercury and Zinc-carbon batteries, are rechargeable. If an external voltage is supplied to the battery it will produce an electrolysis of the products in the overall reaction (discharge), thus recovering initial components which made the battery work.

Solid state Lithium battery

Main article: Lithium battery

Most of the batteries work using an aqueous electrolyte or a moist electrolyte paste instead, however a solid state battery operates using a solid electrolyte. Solid state lithium batteries are an example of this, where a solid Lithium bar acts as the anode, a bar of Lithium sulfide or Vanadium oxide acts as the cathode and a polymer, allowing the passage of ions and not electrons, serves as the electrolyte. The advantage of this kind of battery from others is that Lithium possess the highest negative value of standard reduction potential. It is also a light metal and therefore less mass is required to generate 1 mole of electrons. This battery is rechargeable and it can provide a direct current of about 3 V. Although solid state batteries are frowned upon nowadays, it is likely they will someday become a reliable source of electricity.

Flow battery/ Redox flow battery

Main 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 vanadium redox battery and zinc bromine battery.

Fuel cells

Main article: Fuel cell

Fossil fuels are used on power plants to supply electrical needs of a certain area, however the conversion of them into electricity is a low efficient process, in fact the most efficient electrical power plant it may convert into electricity about 40% of the original chemical energy when burned or processed.

To enhance electrical production, scientists developed fuel cells where combustion reactions are stimulated by electrochemical methods, thus requiring continuous replenishment of the reactants consumed.

The most popular is the oxygen-hydrogen fuel cell, where two inert–electrodes (porous electrodes of Nickel and Nickel oxide) are placed in an electrolytic solution such as hot caustic potash, in both compartments (anode and cathode) gaseous hydrogen and oxygen are bubbled into solution.

Oxygen-hydrogen fuel cell reactions are shown bellow:

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 some-like to hydrogen combustion, differing on oxidation and reduction took place in anode and cathode separately, similar to the electrode used in the cell for measuring standard reduction potential having a double function acting as electrical conductors providing a surface required to decomposition of the molecules into atoms before electron transferring, thus named electrocatalysts. Platinum, nickel, rhodium are good electrocatalysts.

Corrosion

Main article: 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 dollars per year.

Iron corrosion

File:Iron-rusting-scheme.PNG

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:

  1. 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. Electrons are transferred from iron reducing oxygen in the atmosphere into 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; H+ ions come from reaction between carbon dioxide in the atmosphere and water, forming 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)

Iron(III) oxide 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 oxidation, this explains why rusting is quicker on salt water.

Corrosion of coinage metals

Coinage metals, such as copper and silver, can also slowly corrode. At standard temperature and pressure, a patina of green-blue copper carbonate forms on the surface of copper. Silver cutlery that is in contact with food can develop a layer of Silver sulfide.

Prevention of Corrosion

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

While it is almost impossible to prevent anode/cathode formation, if a non-conducting material covers the metal contact with the electrolyte is not possible and corrosion will not occur.

Coating

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

Other prevention is called passivation where a metal is coated with another metal such as tin can. Tin is a metal that rapidly corrodes to form a mono-molecular oxide coating that prevents further corrosion of the tin. The tin prevents the electrolyte from reaching the base metal, usually steel (iron). However, if the tin coating is scratched the iron becomes anodic and the can corrodes rapidly.

Sacrificial anodes

A method commonly used to protect a structural metal is to attach a metal which is more anodic than the metal to be protected. This forces the structural metal to be cathodic, thus spared corrosion. It is called "sacrificial" because the anode dissolves and has to be replaced periodically.

Zinc bars are attached at various locations on steel ship hulls to render the ship hull cathodic. The zinc bars are replaced periodically. Other metals, such as magnesium, would work very well but zinc is the least expensive useful metal.

To protect pipelines, 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.


See also

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External links


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