Corrosion (Application of redox reaction)

Corrosion
Millions of dollars are lost each year because of corrosion. Much of this loss is due to the corrosion of iron and steel, although many other metals may corrode as well. The problem with iron as well as many other metals is that the oxide formed by oxidation does not firmly adhere to the surface of the metal and flakes off easily causing "pitting". Extensive pitting eventually causes structural weakness and disintegration of the metal. (It should be noted, however, that certain metals such as aluminum, form a very tough oxide coating which strongly bonds to the surface of the metal preventing the surface from further exposure to oxygen and corrosion).

Corrosion occurs in the presence of moisture. For example when iron is exposed to moist air, it reacts with oxygen to form rust, 

F2O3.XH2O

The amount of water complexed with the iron (III) oxide (ferric oxide) varies as indicated by the letter "X". The amount of water present also determines the color of rust, which may vary from black to yellow to orange brown. The formation of rust is a very complex process which is thought to begin with the oxidation of iron to ferrous (iron "+2") ions.
  Fe -------> Fe+2 + 2 e-
Both water and oxygen are required for the next sequence of reactions. The iron (+2) ions are further oxidized to form ferric ions (iron "+3") ions.
  Fe+2 ------------> Fe+3 + 1 e-
Tthe electrons provided from both oxidation steps are used to reduce oxygen as shown.
  O2 (g) + 2 H2O + 4e- ------> 4 OH- 
The ferric ions then combine with oxygen to form ferric oxide [iron (III) oxide] which is then hydrated with varying amounts of water. The overall equation for the rust formation may be written as :

4Fe2+(aq)   +  O2(g)  +  [ 4 + 2xH2O (l) ] ---------> 2Fe2O3.XH2O (s)  +  8H+ (aq)

                                                                                              rust

The formation of rust can occur at some distance away from the actual pitting or erosion of iron as illustrated below. This is possible because the electrons produced via the initial oxidation of iron can be conducted through the metal and the iron ions can diffuse through the water layer to another point on the metal surface where oxygen is available. This process results in an electrochemical cell in which iron serves as the anode, oxygen gas as the cathode, and the aqueous solution of ions serving as a "salt bridge" as shown below. 

The involvement of water accounts for the fact that rusting occurs much more rapidly in moist conditions as compared to a dry environment such as a desert. Many other factors affect the rate of corrosion. For example the presence of salt greatly enhances the rusting of metals. This is due to the fact that the dissolved salt increases the conductivity of the aqueous solution formed at the surface of the metal and enhances the rate of electrochemical corrosion. This is one reason why iron or steel tend to corrode much more quickly when exposed to salt (such as that used to melt snow or ice on roads) or moist salty air near the ocean.

Photo-oxidation - Photochromic Glass (Application of redox reaction)

Photo-oxidation - Photochromic Glass
Many people who wear eye glasses prefer those made with photochromic lenses or glass lenses which darken when exposed to bright light. These eyeglasses eliminate the need for sunglasses as they can reduce up to 80% of transmitted light. The basis of this change in color in response to light can be explained in terms of oxidation-reduction reactions. Glass consists of a complex matrix of silicates which is ordinarily transparent to visible light. In photochromic lenses, silver chloride (AgCl) and copper (I) chloride (CuCl) crystals are added during the manufacturing of the glass while it is in the molten state and these crystals become uniformly embedded in the glass as it solidifies. One characteristic of silver chloride is its suscepitibility to oxidation and reduction by light as described below.
  Cl- -----------> Cl + e-
  oxidation
 
  Ag+ + e- -----------> Ag
  reduction
The chloride ions are oxidized to produce chlorine atoms and an electron. The electron is then transferred to silver ions to produce silver atoms. These atoms cluster together and block the transmittance of light, causing the lenses to darken. This process occurs almost instantaneously. As the degree of "darkening" is dependent on the intensity of the light, these photochromic lenses are quite convenient and all but eliminate the need for an extra pair of sunglasses.
The photochromic process would not be useful unless it were reversible. The presence of copper (I) chloride reverses the darkening process in the following way. When the lenses are removed from light, the following reactions occur:
  Cl + Cu+ ------> Cu+2 + Cl- 
oxidizing agent reducing agent oxidized species reduced species
The chlorine atoms formed by the exposure to light are reduced by the copper ions, preventing their escape as gaseous atoms from the matrix. The copper (+1) ion is oxidized to produce copper (+2) ions, which then reacts with the silver atoms as shown.
  Cu+2 + Ag ------> Cu+1 + Ag+
oxidizing agent reducing agent reduced species oxidized species
The net effect of these reactions is that the lenses become transparent again as the silver and chloride atoms are converted to their original oxidized and reduced states.

Electrochemical Cells (Application of redox reaction)

Electrochemical Cells
Many oxidation-reduction reactions occur spontaneously, giving off energy. An example involves the spontaneous reaction that occurs when zinc metal is placed in a solution of copper ions as described by the net ionic equation shown below.
  Cu+2 (aq) + Zn (s) -------> Cu(s) + Zn+2 (aq)
The zinc metal slowly "dissolves" as its oxidation produces zinc ions which enter into solution. At the same time, the copper ions gain electrons and are converted into copper atoms which coats the zinc metal or sediments to the bottom of the container. The energy produced in this reaction is quickly dissipated as heat, but it can be made to do useful work by a device called, an electrochemical cell. This is done in the following way.

An electrochemical cell is composed to two compartments or half-cells, each composed of an electrode dipped in a solution of electrolyte. These half-cells are designed to contain the oxidation half-reaction and reduction half-reaction separately as shown below.

The half-cell, called the anode, is the site at which the oxidation of zinc occurs as shown below.
  Zn (s) ----------> Zn+2 (aq) + 2e-
During the oxidation of zinc, the zinc electrode will slowly dissolve to produce zinc ions (Zn+2), which enter into the solution containing Zn+2 (aq) and SO4-2 (aq) ions.
The half-cell, called the cathode, is the site at which reduction of copper occurs as shown below.
  Cu+2 (aq) + 2e- -------> Cu (s)
When the reduction of copper ions (Cu+2) occurs, copper atoms accumulate on the surface of the solid copper electrode.
The reaction in each half-cell does not occur unless the two half cells are connected to each other.
Recall that in order for oxidation to occur, there must be a corresponding reduction reaction that is linked or "coupled" with it. Moreover, in an isolated oxidation or reduction half-cell, an imbalance of electrical charge would occur, the anode would become more positive as zinc cations are produced, and the cathode would become more negative as copper cations are removed from solution. This problem can be solved by using a "salt bridge" connecting the two cells as shown in the diagram below. A "salt bridge" is a porous barrier which prevents the spontaneous mixing of the aqueous solutions in each compartment, but allows the migration of ions in both directions to maintain electrical neutrality. As the oxidation-reduction reaction occurs, cations ( Zn+2) from the anode migrate via the salt bridge to the cathode, while the anion, (SO4)-2, migrates in the opposite direction to maintain electrical neutrality.
The two half-cells are also connected externally. In this arrangement, electrons provided by the oxidation reaction are forced to travel via an external circuit to the site of the reduction reaction. The fact that the reaction occurs spontaneously once these half cells are connected indicates that there is a difference in potential energy. This difference in potential energy is called an electomotive force (emf) and is measured in terms of volts. The zinc/copper cell has an emf of about 1.1 volts under standard conditions.

Any electrical device can be "spliced" into the external circuit to utilize this potential energy produced by the cell for useful work. Although the energy available from a single cell is relatively small, electrochemical cells can be linked in series to boost their energy output. A common and useful application of this characteristic is the "battery". An example is the lead-acid battery used in automobiles. In the lead-acid battery, each cell has a lead metal anode and lead (IV) oxide (lead dioxide) cathode both of which are immersed in a solution of sulfuric acid. This single electrochemical cell produces about 2 volts. Linking 6 of these cells in series produces the 12-volt battery found in most cars today. One disadvantage of these "wet cells" such as the lead-acid battery is that it is very heavy and bulky. However, like many other "wet cells", the oxidation-reduction reaction which occurs can be readily reversed via an external current such as that provided by an automobile's alternator. This prolongs the lifetime and usefulness of such devices as an energy source.

The "Dry-Cell" Battery

The "Dry-Cell" Battery

The most common type of battery used today is the "dry cell" battery. There are many different types of batteries ranging from the relatively large "flashlight" batteries to the minaturized versions used for wristwatches or calculators. Although they vary widely in composition and form, they all work on the sample principle. A "dry-cell" battery is essentially comprised of a metal electrode or graphite rod (elemental carbon) surrounded by a moist electrolyte paste enclosed in a metal cylinder as shown below.

In the most common type of dry cell battery, the cathode is composed of a form of elemental carbon called graphite, which serves as a solid support for the reduction half-reaction. In an acidic dry cell, the reduction reaction occurs within the moist paste comprised of ammonium chloride (NH4Cl) and manganese dioxide (MnO2):
  2 NH4+ + 2 MnO2 + 2e- ------> Mn2O3 + 2 NH3 + H2O
A thin zinc cylinder serves as the anode and it undergoes oxidation:
  Zn (s) ---------------> Zn+2 + 2e-
This dry cell "couple" produces about 1.5 volts. ( These "dry cells" can also be linked in series to boost the voltage produced). In the alkaline version or "alkaline battery", the ammonium chloride is replaced by KOH or NaOH and the half-cell reactions are:
  Zn + 2 OH- -------> ZnO + H2O + 2e-
  2 MnO2 + 2e- + H2O -------> Mn2O3 + 2 OH-
The alkaline dry cell lasts much longer as the zinc anode corrodes less rapidly under basic conditons than under acidic conditions.
Other types of dry cell batteries are the silver battery in which silver metal serves as an inert cathode to support the reduction of silver oxide (Ag2O) and the oxidation of zinc (anode) in a basic medium. The type of battery commonly used for calculators is the mercury cell. In this type of battery, HgO serves as the oxidizing agent (cathode) in a basic medium, while zinc metal serves as the anode. Another type of battery is the nickel/cadmium battery, in which cadmium metal serves as the anode and nickel oxide serves as the cathode in an alkaline medium. Unlike the other types of dry cells described above, the nickel/cadmium cell can be recharged like the lead-acid battery.