Oxidation and Reduction

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Oxidation and Reduction are the result of an atom or molecule losing or gaining an electron. Before that makes any sense, you need to understand that atoms are made of a “nucleus” containing positively-charged protons and a “cloud” of an equal number of negatively-charged electrons orbiting around it, as shown in the familiar “solar system” icon for atoms and atomic energy. Some of the outermost electrons may be lost rather easily (or not, depending on the element), thus destroying the balance between protons and electrons and changing the character of the atom greatly. When an electron is lost, the protons outnumber the electrons by one + charge, and the atom as a whole then becomes oxidized with a +1 charge, as in the sodium ion, Na+, Conversely, if the electron cloud is capable of accepting another electron, then the electrons would outnumber the protons by one – charge and then the atom would become reduced and have a -1 charge, as in the Chloride ion, Cl. In their elemental forms of sodium metal and chlorine gas, the atoms are balanced and electrically neutral and have a net state) of zero.

Consider the enormous difference between a piece of sodium metal and a cloud of greenish chlorine gas on the one hand and a spoonful of salt crystals or sea water on the other. Yet they are almost identical-both are a collection of sodium and chlorine atoms, and they both have the same number of electrons over all, but in the salt crystals, the sodium atoms all have one electron too few, and the chlorine atoms all have one electron too many. They balance out in the NaCI crystal, but that is not the same. Na+ ion is simply different from Na0, and CI ion is nothing like Cl0. So, electrons are extremely important to chemistry, and when an atom becomes oxidized or reduced, its chemistry changes completely. The iron atoms in the centers of the hemoglobin molecules in our red blood cells must be in the +2 oxidation state-ferrous iron, Fe+2 – to work properly in carrying oxygen to our tissues. If they get oxidized to ferric iron, Fe+3, by nitrite, cyanide, or carbon monoxide, for example, the hemoglobin doesn’t work properly, and we die. One tiny electron on each iron atom makes all the difference.

Oxidation and reduction always occur together, because they are like the two sides of a coin. When an atom or molecule loses an electron (becomes oxidized), that electron must go somewhere, and the atom or molecule that accepts it becomes reduced. When a “reducing agent” does its job, it becomes oxidized in the process. When an “oxidizing agent” does its job, it becomes reduced in the process. Oxidation is loss of electrons, and reduction is acceptance of those same electrons by a “partner” in the reaction. Every chemical has its own unique tendency to gain or lose electrons, called its Oxidation-reduction, or “redox” potential, or ORP, which is expressed in volts and is available from reference books. These voltages can be used to calculate whether a particular reaction will “go” without introducing energy from the outside, but that will not be discussed here. Some examples:

  • The simple battery used for teaching electricity, made by dipping a zinc bar and a copper bar into a solution of copper sulfate: Zinc loses electrons more easily than copper does, so the zinc becomes oxidized and the copper becomes reduced. If you simply dip a zinc bar into a solution of copper sulfate, you get a copper-plated zinc bar which is badly corroded, because every atom of copper that “plates out” on the zinc bar is matched by a zinc ion liberated into the solution. If you use both bars and connect them with a wire to form an electrical circuit, the plating process will proceed until the zinc bar disintegrates and breaks the circuit. The reaction below describes both situations:
    Cu+2 + Zn0 <=> Cu0 + Zn+2
  • The oxidation of hydrogen sulfide (“rotten egg” smell) by chlorine: if only the minimum amount of chlorine is used, the result is elemental sulfur, which is a solid that must be filtered afterward. However, it is possible to oxidize the sulfur atom all the way to sulfate if more chlorine is used. The first reaction below shows the sulfur being oxidized from the -2 state in H2S to the zero state in elemental S, and the second one shows the sulfur being oxidized to the +6 state as sulfuric acid.
    H2S + CI2 <=> 2H+ + 2CI + S
    H2S + 4CI2 + 4H2O <=> 8H+ + 8CI + H2SO4

Most elements have at least two possible oxidation states. Note chlorine, below, which has more possibilities than most elements:

Chemical Oxidation State Name Comments
Perchlorate ion
Chlorate ion
Chlorine dioxide
Chlorite ion
Hypochlorite ion
Chlorine gas
Chloride ion
Rocket fuel
Gas used for oxidation and disinfection
Bleaching agent
Laundry bleach
Elemental chlorine
Table salt, sea water

This also demonstrates some useful rules of nomenclature:

“-ate” = suffix denoting a “high” oxidation state; same as “-ic” suffix for acids*
“-ite” = suffix denoting a “low” oxidation state; same as “-ous” suffix for acids*
“per-” (short for “hyper-”) = prefix denoting “higher oxidation state than ‘-ate’”
“hypo-” = prefix denoting “lower oxidation state than ‘-ite’”
“-ide” = suffix denoting the complete absence of oxygen

*names of some acids: nitric acid, nitrous acid, sulfuric acid, sulfurous acid, perchloric acid, chloric acid, chlorous acid, hypochlorous acid.

The Electromotive Series is a listing of metallic elements in descending order of their oxidation-reduction potential or ORP. One of the characteristics of metals is the free movement of the outer “valence” electrons, which makes them good electrical, conductors. The ORP voltage is a measure of each element’s readiness to lose electrons (be oxidized). They are arranged at right, with the most reactive element at the top and the least reactive at the bottom. Elemental potassium and sodium are so reactive that they react with water as if it were strong acid, making NaOH or KOH and hydrogen gas. At the other end, everyone knows that silver, platinum and gold are “noble” metals that are difficult or impossible to corrode. This list can also be used in the reverse: the ions of a metal can cause the corrosion of any metal above it in the list. For example, if a copper wire is immersed in a solution of silver ions, the Ag+ -> Ag (plating) and the Cu -> Cu+ (corrosion). Thus, silver metal is relatively non-corrodable, but silver ion is highly corrosive. Alchemists’ old name for silver nitrate (AgNO3) was “lunar caustic” – the Ag+ ion is so caustic and corrosive that it will even react with the proteins in your flesh. Perhaps that is why Ag+ and Cu++ have some weak antibacterial activity, even at low concentrations.