Oxidation-reduction

Originally chemists viewed oxidation as a class of chemical reactions in which a chemical species (e.g., a chemical element or compound) reacts with oxygen to form an oxygen-containing product referred to as an 'oxide'. In modern terminology, chemists would describe that as the chemical species reacting with an oxygen molecule (O₂) such that it combines with an atom of oxygen (O) to form an oxygen-containing product, as when hydrogen (H₂) reacts with O₂ to form the oxygen-containing product, H₂O, namely water:

2H₂ + O₂ → 2H₂O [or, 2HOH]

In that reaction, each molecular pair of hydrogen atoms has been 'oxidized' by gaining an oxygen atom.

History

The 'Father of Chemistry', Antoine-Laurent Lavoisier (1743-1794), in his Elements of Chemistry (originally published in French in 1789), writes of oxidation in relation to the products formed when metals (e.g., mercury, iron) are exposed to air and a certain amount of heat:

The term oxidation, or calcination, is chiefly used to signify the process by which metals exposed to a certain degree of heat are converted to oxides, by absorbing oxygen from the air.

Lavoisier recognized calcination of metals as oxidation, the resulting oxides referred to as calxes, or calces, (sing., (calx).

Lavoisier's work revealed the common link — namely the requirement for the oxygen component of air, the chemical reaction of oxidation — among the air-requiring processes of:

• burning (combustion),

• rusting and other such so-called calcinations of metals, and

• breathing (respiration) by animals .

Prior to Lavoisier, the chemist, Georg Ernst Stahl (1660-1734), taught that oxidation, as it came to be called later, involved the escape of a common constituent of combustible substances, including metals capable of calcination, referred to as phlogiston, whose concentration differed among combustible/calcinable substances. Coal, which burns violently, he considered nearly all phlogiston.

The phenomenon of converting a calx back to the metal it started as, by reacting the calx with burning coal, Stahl explained as returning the phlogiston back to the metal. That phenomenon was known as 'reduction', i.e., reducing a larger amount of calx to the smaller amount of the pure metal, the weight difference contrary to Stahl's 'phlogiston theory'.

As the concept of oxidation matured with the development of the atomic theory, the concept of reduction also became refined, and more importantly, became inextricably linked to oxidation as a simultaneous event that accompanied every oxidation event (see below). For example, in the oxidation (oxygenation) of dihydrogen with dioxygen to produce water, as described above, the hydrogen atoms were oxidized and, at the same time, the oxygens atoms became reduced.

How Lavoisier came to replace the the phlogiston theory with the oxygenation theory is one of the great stories in the annals of the history of chemistry.

The modern concept of oxidation and reduction

With the advent of the atomic and quantum theories, as chemists continued to study the reaction of metals with oxygen, they learned that such reactions involved a transfer of electrons, specifically a transfer of electrons from the metal to oxygen. For example, when the elemental metal, magnesium, Mg, a solid under ordinary conditions, reacts with molecular oxygen, O₂, present in air, magnesium atoms exposed to oxygen each transfer two electrons, 2e^-, to an oxygen atom, forming the ionic compound, magnesium oxide, MgO. Many such magnesium oxide compounds together form a crystalline structure. The balanced 'molecular' chemical reaction is written:

2Mg(s) + O₂(g) → 2MgO(s), where s=solid, g=gas

A visible sample of elemental magnesium exposed to air acquires a sheen of magnesium oxide crystals that protects the underlying magnesium ions from the oxidation reaction.

Why does the electron transfer occur? A precise answer requires an understanding of quantum mechanics, which reveals that oxygen atoms have a much greater 'affinity' for electrons than do magnesium atoms. Metaphorically, and oversimplified, atoms in their neutral state behave 'as if' they wanted to fill to its natural capacity their so-called valence electron shell. An oxygen atom, a stronger electron attractor than a magnesium atom, requires two electrons to accomplish that goal.

Some definitions

At this point we introduce a few helpful definitions:

— Reaction Type — — Definition — — Representative Examples — — Comments —   OXIDATION   -- loss of electrons -- (oxidation = 'de-electronation') K → K+ + e- O2•- → O2 + e-   oxidation of a potassium atom to a potassium ion oxidation of a superoxide radical to oxygen   OXIDATION -- loss of electron with hydrogen nucleus -- --   examples depend on the atom or molecule that accepts the hydrogen atom   OXIDATION -- gain of oxygen -- C + O2 → CO2   oxidation of carbon to carbon dioxide (the more electronegative oxygen obtains greater share of carbon's electrons)   REDUCTION -- gain of electrons -- (reduction = 're-electronation') Cl + e- → Cl- O2 + e- → O2•-   reduction of a chlorine atom to a chloride ion reduction of an oxygen molecule to a superoxide radical   REDUCTION -- gain of electron with a hydrogen nucleus -- C +2H2 → CH4   reduction of a carbon atom to a methane molecule (the carbon atom gains electrons it shares with a hydrogen nucleus)   REDUCTION -- loss of oxygen -- CO2 + C → 2CO   reduction of carbon dioxide to carbon monoxide concomitant oxidation of carbon to carbon monoxide