Oxidation - Definition


Oxidation - Definition

The most fundamental reactions in chemistry are the redox processes. The term redox process accounts for all processes in which atoms have their oxidation number changed. This can be a simple redox process, such as the combustion of carbon by oxygen to yield carbon dioxide, it could be the reduction of carbon by hydrogen to yield methane, or it could be the oxidation of sugar in the human body, through a series of very complex electron transfer processes, to yield water and carbon dioxide.

The term redox comes from the two concepts of reduction and oxidation. Reduction descibes the uptake of an electron by a molecule or atom. Oxidation describes the loss of an electron by a molecule or atom. These two terms go together, because in a chemical reaction, one cannot occur without the other; electrons lost by one compound must be gained by another.

Contents

1 Oxidizing and Reducing agents

2 Former meaning (Oxygen/Hydrogen)

3 Examples of redox reactions

4 Redox reactions in biology

5 See also

Oxidizing and Reducing agents

Substances that have the ability to oxidize other substances are said to be oxidative and are known as oxidizing agents/oxidants or oxidizers. Put in another way, the oxidant removes electrons from the substance. Oxidants are usually chemical substances in high oxidation numbers (eg. MnO₄^-, CrO₃, OsO₄) or very electronegative substances that can gain one or two extra electrons by oxidizing a substance (O₂, O₃, F₂, Cl₂, Br₂)

Substances that have the ability to reduce other substances are said to be reductive and are known as reductive agents/reductants or reducers. Put in another way, the reductant transfers electrons to the substance. Reductants in chemistry are very diverse. Metal reduction - electropositive elemental metals can be used (Li, Na, Mg, Fe, Zn, Al). These metals are very eager to give away electrons. Other kinds of reductants are hydride transfer reagents (NaBH₄, LiAlH₄), these reagents are widely used in organic chemistry, primarily in the reduction of carbonyl compounds to alkohols. Another useful method is reductions involving hydrogen gas (H₂) with a palladium, platinium or nickel catalyst. These catalytic reductions are primarily used in the reduction of carbon-carbon double or triple bonds.

The chemical way to look at redox processes is that the reductant transfers electrons to the oxidant. Thus, at the end of the reaction, the reductant will have been oxidized and the oxidant will have been reduced. This does not mean however, that the reverse process takes place (because that would lead to status quo), but it does often lead to equlibrium.

Chemistry students sometimes remember the nature of oxidation and reduction using the simple mnemonic device "LEO says GER" - the Loss of Electrons is Oxidation, while the Gain of Electrons is Reduction. Another is "OIL RIG" - Oxidation Is Loss, Reduction Is Gain.

Former meaning (Oxygen/Hydrogen)

Formerly, oxidation (reduction) simply meant the addition (removing) of oxygen or the removing (addition) of hydrogen (hence the name oxidation). However, when the terms are now used it is normally in the more general sense.

Examples of redox reactions

A good example is the reaction between hydrogen and fluorine:

H₂ + F₂ → 2HF

We can write this overall reaction as two half-reactions: an oxidation reaction:

H₂ → 2H⁺ + 2e^-

and a reduction reaction:

F₂ + 2e^- → 2F^-

Elements always have an oxidation number of zero. In the first half reaction hydrogen is oxidized from an oxidation number of zero to an oxidation number of +1. In the second half reaction fluorine is reduced from an oxidation number of zero to an oxidation number of −1.

When adding the reactions together the electrons cancel:

        H₂ → 2H⁺ + 2e^-
+ 2e^- + F₂ → 2F^-
  ---------------------
   H₂ + F₂ → 2H⁺ + 2F^-

And the ions combine to form hydrogen fluoride:

2H⁺ + 2F^- → 2HF

Here are some other examples:

iron(II) oxidises/oxidizes to iron(III):

Fe²⁺ → Fe³⁺ + e^-

hydrogen peroxide reduces to hydroxide:

H₂O₂ + 2 e^- → 2 OH^-

overall equation for the above:

2Fe²⁺ + H₂O₂ + 2H⁺ → 2Fe³⁺ + 2H₂O

iron(II) oxidises/oxidizes to and oxygen reduces to iron(III) oxide (commonly known as rusting or tarnishing):

4Fe + 3O₂ → 2 Fe₂O₃.

Another example is the burning of hydrocarbons to produce water, carbon dioxide, some partially oxidized forms, and heat energy. Complete oxidation of materials containing carbon produces carbon dioxide, which is linked to global warming because it absorbs certain wavelengths of infrared light.

In organic chemistry, stepwise oxidation of a hydrocarbon produces water and, successively, an alcohol, an aldehyde or a ketone, carboxylic acid, and then a peroxide.

In inorganic chemistry terms, incompletely oxidized carbon takes the form of carbonate, bicarbonate or carbon monoxide.

Redox reactions in biology

Much biological energy is stored and released by means of redox reactions. Photosynthesis involves the reduction of carbon dioxide into sugars and the oxidation of water into molecular oxygen. The reverse reaction, respiration, oxidises/oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce NAD⁺, which then contributes to the creation of a proton gradient, which drives the synthesis of ATP and is maintained by the reduction of oxygen.

In animal cells, mitochondria perform similar functions.

The term redox state is often used to describe the balance of NAD⁺/NADH and NADP⁺/NADPH in a biological system such as a cell or organ. The redox state is reflected in the balance of several sets of metabolites (e.g., lactate and pyruvate, beta-hydroxybutyrate and acetoacetate) whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as hypoxia, shock, and sepsis.