Electrochemistry is the study of chemical
reactions which take place at the interface of an electrode,
usually a solid metal
or a semiconductor,
and an ionic conductor, the electrolyte. These reactions involve electric
charges moving between the electrodes and the electrolyte (or ionic species in a solution). Thus
electrochemistry deals with the interaction between electrical energy and
chemical change.
When a chemical reaction is caused by an externally supplied
current,
as in electrolysis,
or if an electrical current is produced by a spontaneous chemical reaction as
in a battery, it is called an electrochemical
reaction. Chemical reactions where electrons are
transferred directly between molecules and/or atoms are called oxidation-reduction or (redox) reactions. In general,
electrochemistry describes the overall reactions when individual redox reactions are
separate but connected by an external electric circuit and an intervening
electrolyte.
History
Understanding of electrical matters began in the sixteenth
century. During this 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—18th century the French chemist Charles François de Cisternay du
Fay had 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.
In 1785, Charles-Augustin de Coulomb developed
the law of electrostatic attraction as an outgrowth of his attempt
to investigate the law of electrical repulsions as stated by Joseph
Priestley in England.
In the late 18th century 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.
In 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, William Nicholson and Johann Wilhelm Ritter succeeded in decomposing
water into hydrogen
and oxygen by electrolysis.
Soon thereafter Ritter discovered the process of electroplating.
He also observed that the amount of metal deposited and the amount of oxygen
produced during an electrolytic process depended on the distance between the electrodes.
By 1801, Ritter observed thermoelectric
currents and anticipated the discovery of thermoelectricity by Thomas Johann Seebeck.
By the 1810s, William Hyde Wollaston made improvements to
the galvanic
cell. 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.
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 led him to state his two laws of electrochemistry. In
1836, John Daniell invented a primary cell which
solved the problem of polarization by eliminating hydrogen gas
generation at the positive electrode. Later results revealed that alloying the amalgamated zinc with mercury would produce a higher voltage.
William Grove produced the first fuel cell
in 1839. In 1846, Wilhelm Weber developed the electrodynamometer. In 1868, Georges Leclanché patented a new cell which
eventually became the forerunner to the world's first widely used battery, the zinc carbon cell.
Svante 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 an efficient method
(the Hall–Héroult process) to obtain aluminium
using electrolysis of molten alumina.
In 1894, Friedrich
Ostwald concluded important studies of the conductivity and electrolytic dissociation
of organic
acids.
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.
20th century and recent developments
In 1902, The Electrochemical Society (ECS) was
founded.
In 1909, Robert Andrews Millikan began a series of
experiments (see oil drop experiment) to determine the electric
charge carried by a single electron.
In 1923, Johannes Nicolaus Brønsted and Martin
Lowry published essentially the same theory about how acids and bases
behave, using an electrochemical basis.
In 1937, Arne Tiselius developed the first sophisticated electrophoretic
apparatus. 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 quantum electrochemistry was developed by Revaz
Dogonadze and his pupils.
Principles
Oxidation and reduction
The term "redox" stands for reduction-oxidation. It refers to
electrochemical processes involving electron
transfer to or from a molecule or ion changing its oxidation
state. This reaction can occur through the application of an external voltage or
through the release of chemical energy. Oxidation and reduction describe the
change of oxidation state that takes place in the atoms, ions or molecules
involved in an electrochemical reaction. Formally, oxidation state is the
hypothetical charge that an atom would have if all bonds to
atoms of different elements were 100% ionic. An
atom or ion that gives up an electron to another atom or ion has its oxidation
state increase, and the recipient of the negatively charged electron has its
oxidation state decrease.
For example, when atomic sodium reacts with
atomic chlorine,
sodium donates one electron and attains an oxidation state of +1. Chlorine
accepts the electron and its oxidation state is reduced to −1. The sign of the
oxidation state (positive/negative) actually corresponds to the value of each
ion's electronic charge. The attraction of the differently charged sodium and
chlorine ions is the reason they then form an ionic bond.
The loss of electrons from an atom or molecule 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"
says "GER" (Lose Electrons: Oxidation, Gain Electrons:
Reduction). Oxidation and reduction always occur in a paired fashion such that
one species is oxidized when another is reduced. 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. Thus,
the oxidizing agent is always being reduced in a reaction; the reducing agent
is always being oxidized. Oxygen is a common oxidizing agent, but not the only
one. Despite the name, an oxidation reaction does not necessarily need to
involve oxygen. In fact, a fire can be fed by an oxidant other than oxygen; fluorine fires
are often unquenchable, as fluorine is an even stronger oxidant (it has a
higher electronegativity and thus accepts electrons even
better) than oxygen.
For reactions involving oxygen, the gain of oxygen implies
the oxidation of the atom or molecule to which the oxygen is added (and the
oxygen is reduced). In organic compounds, such as butane or ethanol, the loss
of hydrogen implies oxidation of the molecule from which it is lost (and the
hydrogen is reduced). This follows because the hydrogen donates its electron in
covalent bonds with non-metals but it takes the electron along when it is lost.
Conversely, loss of oxygen or gain of hydrogen implies reduction.
Balancing redox reactions
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.
Acidic 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.
Unbalanced reaction: Mn2+(aq) +
NaBiO3(s) → Bi3+(aq) + MnO4–(aq)
Oxidation: 4 H2O(l) + Mn2+(aq)
→ MnO4–(aq) + 8 H+(aq) +
5 e–
Reduction: 2 e– + 6 H+(aq)
+ BiO3–(s) → Bi3+(aq) +
3 H2O(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.
8 H2O(l) + 2 Mn2+(aq)
→ 2 MnO4–(aq) + 16 H+(aq)
+ 10 e–
10 e– + 30 H+(aq) + 5 BiO3–(s)
→ 5 Bi3+(aq) + 15 H2O(l)
Reaction balanced:
14 H+(aq) + 2 Mn2+(aq)
+ 5 NaBiO3(s) → 7 H2O(l) + 2 MnO4–(aq)
+ 5 Bi3+(aq) + 5 Na+(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.
Unbalanced reaction: KMnO4 + Na2SO3
+ H2O → MnO2 + Na2SO4 + KOH
Reduction: 3 e– + 2 H2O + MnO4–
→ MnO2 + 4 OH–
Oxidation: 2 OH– + SO32–
→ SO42– + H2O + 2 e–
The same procedure as followed on acid medium by multiplying
electrons to opposite half reactions solve the equation thus balancing the overall
reaction.
6 e– + 4 H2O + 2 MnO4–
→ 2 MnO2 + 8 OH–
6 OH– + 3 SO32– → 3 SO42–
+ 3 H2O + 6e–
Equation balanced:
2 KMnO4 + 3 Na2SO3 + H2O
→ 2 MnO2 + 3 Na2SO4 + 2 KOH
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.
Unbalanced reaction: C3H8 + O2
→ CO2 + H2O
Reduction: 4 H+ + O2 + 4 e–
→ 2 H2O
Oxidation: 6 H2O + C3H8
→ 3 CO2 + 20 e– + 20 H+
As in acid and basic medium, electrons which were used to
compensate oxidation changes are multiplied to opposite half reactions, thus
solving the equation.
20 H+ + 5 O2 + 20 e– → 10 H2O
6 H2O + C3H8 → 3 CO2
+ 20 e– + 20 H+
Equation balanced:
C3H8 + 5 O2 → 3 CO2
+ 4 H2O
Electrochemical cells
An electrochemical cell is a device that produces an
electric current from energy released by a spontaneous redox reaction. This kind of cell
includes the Galvanic cell or Voltaic cell, named after Luigi
Galvani and Alessandro Volta, both scientists who conducted
several experiments on chemical reactions and electric current during the late
18th century.
Electrochemical cells have two conductive electrodes (the
anode and the cathode). The anode is defined as the electrode where oxidation occurs and
the cathode is
the electrode where the reduction takes place. Electrodes can be made from any
sufficiently conductive materials, such as metals, semiconductors, graphite,
and even conductive polymers. In between these electrodes
is the electrolyte,
which contains ions that can freely move.
The galvanic cell uses two different metal electrodes, each
in an electrolyte where the positively charged ions are the oxidized form of
the electrode metal. One electrode will undergo oxidation (the anode) and the
other will undergo reduction (the cathode). The metal of the anode will
oxidize, going from an oxidation state of 0 (in the solid form) to a positive
oxidation state and become an ion. At the cathode, the metal ion in solution
will accept one or more electrons from the cathode and the ion's oxidation state
is reduced to 0. This forms a solid metal that electrodeposits
on the cathode. The two electrodes must be electrically connected to each
other, allowing for a flow of electrons that leave the metal of the anode and
flow through this connection to the ions at the surface of the cathode. This
flow of electrons is an electrical current that can be used to do work, such as
turn a motor or power a light.
A galvanic cell whose electrodes
are zinc and copper submerged in
zinc
sulfate and copper sulfate, respectively, is known as a Daniell
cell.
Half reactions for a Daniell cell are these:
Zinc electrode (anode): Zn(s) → Zn2+(aq)
+ 2 e–
Copper electrode (cathode): Cu2+(aq) +
2 e– → 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 shot glass-shaped container is aerated
with a noble gas and sealed with the Teflon block.
In this example, the anode is zinc metal which oxidizes
(loses electrons) to form zinc ions in solution, and copper ions accept
electrons from the copper metal electrode and the ions deposit at the copper
cathode as an electrodeposit. This cell forms a simple battery as it will spontaneously
generate a flow of electrical current from the anode to the cathode through the
external connection. This reaction can be driven in reverse by applying a
voltage, resulting in the deposition of zinc metal at the anode and formation
of copper ions at the cathode.
To provide a complete electric circuit, there must also be
an ionic conduction path between the anode and cathode electrolytes in addition
to the electron conduction path. The simplest ionic conduction path is to
provide a liquid junction. To avoid mixing between the two electrolytes, the
liquid junction can be provided through a porous plug that allows ion flow
while reducing electrolyte mixing. To further minimize mixing of the
electrolytes, a salt bridge can be used which consists of an
electrolyte saturated gel in an inverted U-tube. As the negatively charged
electrons flow in one direction around this circuit, the positively charged
metal ions flow in the opposite direction in the electrolyte.
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
Daniell cell:
Zn(s) | Zn2+ (1M) || Cu2+
(1M) | 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
To allow prediction of the cell potential, tabulations of standard electrode potential are
available. Such tabulations are referenced to the standard hydrogen electrode
(SHE). The standard hydrogen electrode undergoes
the reaction
2 H+(aq) + 2 e– → H2
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 −0.76 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.
E°cell = E°red (cathode) – E°red
(anode) = E°red (cathode) + E°oxi (anode)
For example, the standard electrode potential for a copper
electrode is:
Cell diagram
Pt(s) | H2 (1 atm) | H+ (1
M) || Cu2+ (1 M) | Cu(s)
E°cell = E°red (cathode) – E°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
Ecell = E°(Cu2+/Cu) – E°(H+/H2)
Or,
E°(Cu2+/Cu) = 0.34 V
Changes in the stoichiometric coefficients of a
balanced cell equation will not change E°red value because the
standard electrode potential is an intensive property.
Spontaneity of redox reaction
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 electric charge transferred through the external
circuit.
Electrical energy = EcellCtrans
where Ecell is the cell potential measured in
volts (V) and Ctrans is the cell current integrated over time and
measured in coulombs (C); Ctrans can also be determined by
multiplying the total number of electrons transferred (measured in moles) times
Faraday's constant (F).
The emf of the cell at zero current is the maximum possible
emf. It is used to calculate the maximum possible electrical energy that could
be obtained from a chemical reaction. This energy is referred to as electrical
work and is expressed by the following equation:
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:
A positive cell potential gives a negative change in Gibbs
free energy. This is consistent with the cell production of an electric
current 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 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 the equilibrium constant, K, and the Gibbs
free energy for an electrochemical cell is expressed as follows:
.
Rearranging to express the relation between standard potential
and equilibrium constant yields
.
The previous equation can use Briggsian logarithm as shown below:
Cell emf dependency on changes in concentration
Nernst equation
The standard potential of an electrochemical cell requires
standard conditions (ΔG°) for all of the reactants. When reactant
concentrations differ from standard conditions, the cell potential will deviate
from the standard potential. In the 20th century German chemist Walther
Nernst proposed a mathematical model to determine the effect of reactant
concentration on electrochemical cell potential.
In the late 19th century, Josiah Willard Gibbs had formulated a theory
to predict whether a chemical reaction is spontaneous based on the free energy
Here ΔG is change in Gibbs
free energy, ΔG° is the cell potential when Q is equal to 1, T
is absolute temperature(Kelvin), R is the gas
constant and Q is reaction
quotient which can be found by dividing products by reactants using only
those products and reactants that are aqueous or gaseous.
Gibbs' key contribution was to formalize the understanding
of the effect of reactant concentration on spontaneity.
Based on Gibbs' work, Nernst extended the theory to include
the contribution from electric potential on charged species. As shown in the
previous section, the change in Gibbs free energy for an electrochemical cell
can be related to the cell potential. Thus, Gibbs' theory becomes
Here n is the number of electrons/mole
product, F is the Faraday constant (coulombs/mole),
and ΔE is cell potential.
Finally, Nernst divided through by the amount of charge
transferred to arrive at a new equation which now bears his name:
Assuming standard conditions (T = 25 °C) and R = 8.3145 J/(K·mol), the equation above can
be expressed on base—10 logarithm as shown below:
Concentration cells
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.
An example is an electrochemical cell, where two copper
electrodes are submerged in two copper(II) sulfate solutions, 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 can undergo the same chemistry (although the reaction proceeds in reverse
at the anode)
Cu2+(aq) + 2 e– → Cu(s)
Le Chatelier's principle indicates that
the reaction is more favorable to reduction as the concentration of Cu2+
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) | Cu2+ (0.05 M) || Cu2+
(2.0 M) | Cu(s)
Where the half cell reactions for oxidation and reduction
are:
Oxidation: Cu(s) → Cu2+ (0.05 M) + 2 e–
Reduction: Cu2+ (2.0 M) + 2 e– → Cu(s)
Overall reaction: Cu2+ (2.0 M) → Cu2+
(0.05 M)
The cell's emf is calculated through Nernst equation as
follows:
The value of E° in 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:
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
Many types of battery have been commercialized and represent
an important practical application of electrochemistry. Early wet cells
powered the first telegraph and telephone
systems, and were the source of current for electroplating.
The zinc-manganese dioxide dry cell was the first portable, non-spillable battery
type that made flashlights and other portable devices practical. The mercury
battery using zinc and mercuric oxide provided higher levels of power and
capacity than the original dry cell for early electronic devices, but has been
phased out of common use due to the danger of mercury pollution from discarded
cells.
The lead–acid battery was the first practical
secondary (rechargeable) battery that could have its capacity replenished from
an external source. The electrochemical reaction that produced current was (to
a useful degree) reversible, allowing electrical energy and chemical energy to
be interchanged as needed. Common lead acid batteries contain a mixture of acid
and water, as well as lead plates. The most common mixture used today is 30%
acid. One problem however is if left uncharged acid will crystallize within the
lead plates of the battery rendering it useless. These batteries last an
average of 3 years with daily use however it is not unheard of for a lead acid
battery to still be functional after 7-10 years. Lead-acid cells continue to be
widely used in automobiles.
All the preceding types have water-based electrolytes, which
limits the maximum voltage per cell. The freezing of water limits low
temperature performance. The lithium
battery, which does not (and cannot) use water in the electrolyte, provides
improved performance over other types; a rechargeable lithium ion battery is an essential part of
many mobile devices.
The flow battery, an experimental type, offers the option
of vastly larger energy capacity because its reactants can be replenished from
external reservoirs. The fuel cell can turn the chemical energy bound in
hydrocarbon gases or hydrogen directly into electrical energy with much higher
efficiency than any combustion process; such devices have powered many
spacecraft and are being applied to grid energy storage for the public power
system.
Corrosion
Corrosion is the term applied to steel 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
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)
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.
Fe(s) → Fe2+(aq) + 2 e–
Electrons are transferred from iron reducing oxygen in
the atmosphere
into water on the cathode, which is placed in another
region of the metal.
O2(g) + 4 H+(aq) + 4 e–
→ 2 H2O(l)
Global reaction for the process:
2 Fe(s) + O2(g) + 4 H+(aq)
→ 2 Fe2+(aq) + 2 H2O(l)
Standard emf for iron rusting:
E° = E° (cathode) – E° (anode)
E° = 1.23V – (−0.44 V) = 1.67 V
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:
4 Fe2+(aq) + O2(g) + (4+2x)
H2O(l) → 2 Fe2O3·x H2O +
8 H+(aq)
Iron(III) oxide hydrated is
known as rust. The concentration of water associated with iron oxide varies,
thus chemical representation is presented as Fe2O3·x H2O.
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 common metals
Coinage metals, such as copper and silver, slowly
corrode through use. A patina of green-blue copper
carbonate forms on the surface of copper with
exposure to the water and carbon dioxide in the air. Silver coins or cutlery that are
exposed to high sulfur foods such as eggs
or the low levels of sulfur species in the air develop a layer of black Silver
sulfide.
Gold
and platinum
are extremely difficult to oxidize under normal circumstances, and require
exposure to a powerful chemical oxidizing agent such as aqua regia.
Some common metals oxidize extremely rapidly in air. Titanium and
aluminium oxidize instantaneously in contact with the oxygen in the air. These
metals form an extremely thin layer of oxidized metal on the surface. This thin
layer of oxide protects the underlying layers of the metal from the air
preventing the entire metal from oxidizing. These metals are used in
applications where corrosion resistance is important. Iron, in contrast, has
an oxide that forms in air and water, called rust, that does not
stop the further oxidation of the iron. Thus iron left exposed to air and water
will continue to rust until all of the iron is oxided.
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 can be coated with paint or other less
conductive metals (passivation). This prevents the
metal surface from being exposed to electrolytes.
Scratches exposing the metal substrate will result in corrosion. The region
under the coating adjacent to the scratch acts as the anode of the
reaction.
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 to 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, an ingot of buried or exposed
magnesium (or zinc) is buried
beside the pipeline and is connected electrically
to the pipe above ground. The pipeline is forced to be a cathode and is
protected from being oxidized and rusting. The magnesium anode is sacrificed.
At intervals new ingots
are buried to replace those lost.
Electrolysis
The spontaneous redox reactions of a conventional battery
produce electricity through the different chemical potentials of the cathode
and anode in the electrolyte. However, electrolysis requires an external source
of electrical energy to induce a chemical reaction,
and this process takes place in a compartment called an electrolytic
cell.
Electrolysis of molten sodium chloride
When molten, the salt sodium
chloride can be electrolyzed to yield metallic sodium and gaseous chlorine.
Industrially this process takes place in a special cell named Down's cell. The
cell is connected to an electrical power supply, allowing electrons to
migrate from the power supply to the electrolytic cell.
Reactions that take place at Down's cell are the following:
Anode (oxidation): 2 Cl– → Cl2(g) + 2
e–
Cathode (reduction): 2 Na+(l) + 2 e–
→ 2 Na(l)
Overall reaction: 2 Na+ + 2 Cl–(l)
→ 2 Na(l) + Cl2(g)
This process can yield large amounts of metallic sodium and
gaseous chlorine, and is widely used on mineral
dressing and metallurgy industries.
The emf for this process is approximately −4 V indicating a (very)
non-spontaneous process. In order for this reaction to occur the power supply
should provide at least a potential of 4 V. However, larger voltages must be
used for this reaction to occur at a high rate.
Electrolysis of water
Water can be converted to its component elemental gasses, H2
and O2 through the application of an external voltage. Water doesn't
decompose into hydrogen
and oxygen spontaneously as the Gibbs
free energy for the process at standard conditions is about 474.4 kJ. The
decomposition of water into hydrogen and oxygen can be performed in an
electrolytic cell. In it, a pair of inert electrodes
usually made of platinum
immersed in water act as anode and cathode in the electrolytic process. The
electrolysis starts with the application of an external voltage between the
electrodes. This process will not occur except at extremely high voltages
without an electrolyte such as sodium
chloride or sulfuric acid (most used 0.1 M).
Bubbles from the gases will be seen near both electrodes.
The following half reactions describe the process mentioned above:
Anode (oxidation): 2 H2O(l) → O2(g)
+ 4 H+(aq) + 4 e–
Cathode (reduction): 2 H2O(g) + 2 e–
→ H2(g) + 2 OH–(aq)
Overall reaction: 2 H2O(l) → 2 H2(g)
+ O2(g)
Although strong acids may be used in the apparatus, the
reaction will not net consume the acid. While this reaction will work at any
conductive electrode at a sufficiently large potential, platinum catalyzes
both hydrogen and oxygen formation, allowing for relatively mild voltages (~2 V
depending on the pH).
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, cation,
which is the positive ion, will be attracted to the cathode (-), thus reducing the
sodium ion. The anion will then be
attracted to the anode (+) oxidizing chloride ion.
The following half reactions describes the process
mentioned:
1. Cathode: Na+(aq) + e– →
Na(s) E°red = –2.71 V
2. Anode: 2 Cl–(aq) → Cl2(g)
+ 2 e– E°red = +1.36 V
3. Cathode: 2 H2O(l) + 2 e–
→ H2(g) + 2 OH–(aq) E°red
= –0.83 V
4. Anode: 2 H2O(l) → O2(g)
+ 4 H+(aq) + 4 e– E°red
= +1.23 V
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 and 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 show 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°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:
Anode (oxidation): 2 Cl–(aq) → Cl2(g)
+ 2 e–
Cathode (reduction): 2 H2O(l) + 2 e–
→ H2(g) + 2 OH–(aq)
Overall reaction: 2 H2O + 2 Cl–(aq)
→ H2(g) + Cl2(g) + 2 OH–(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 and Faraday's laws
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.
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 is a simplified equation of Faraday's first law:
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
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 words, 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.
Applications
There are various extremely important electrochemical
processes in both nature and industry, like the coating of objects with metals
or metal oxides through electrodeposition and the detection of alcohol in
drunken drivers through the redox reaction of ethanol. The generation of
chemical energy through photosynthesis is inherently an electrochemical
process, as is production of metals like aluminum and titanium from their ores.
Certain diabetes blood sugar meters measure the amount of glucose in the blood
through its redox potential.
The action potentials that travel down neurons are based
on electric current generated by the movement of sodium and potassium ions into
and out of cells. Specialized cells in certain animals like the electric
eel can generate electric currents powerful enough to disable much larger
animals.
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