As stated above, for the purposes of oxidation and reduction the oxidation number can be thought of as the apparent ionic charge of an atom within a compound. For example, in sulphuric acid the sulphur is in the VI (6+) oxidation state. For the purposes of redox we can consider that it has an ionic charge of +6 (even though it is clearly covalently bonded). This makes it easier to follow any transfer of electrons.
What happens in the clip from Daylight is more violent still: what we know as an explosion is a very rapid, practically instantaneous oxidation, with intense heat and light release, plus an increased pressure due to rapid release of gases (products of the redox reaction).
Hexavalent chromium [Cr(VI)] is a well-recognized human carcinogen that is found in the workplaces of several millions of workers worldwide . Environmental exposure to Cr(VI) has raised questions about its potential adverse health effects in the general population (). Intracellular reduction of Cr(VI) to stable Cr(III) is responsible for the production of DNA-damaging products. Cr(VI) metabolism can generate variable amounts of Cr(V) and Cr(IV) intermediates depending upon the nature of the reducing agent. Ascorbate (Asc) is the key reducer of Cr(VI) in cells in vivo, accounting for more than 90% of its metabolism (, ; ). The reduction of Cr(VI) by Asc generates Cr(IV) as the main intermediate under physiological conditions (; ). Cr(VI) reactions with thiol-based secondary reducing agents yield reactive Cr(V) as the first intermediate (; ). The final product of Cr(VI) reduction by all reducing agents is Cr(III), which forms several mutagenic Cr–DNA adducts (). Human and nonhepatic rodent cells in standard cultures contain either undetectable or low micromolar concentrations of Asc (, ) in contrast to the 1–3 millimolar amounts of Asc in the main tissues in vivo (; ). Consequently, metabolism of Cr(VI) in cultured cells is dominated by the most abundant thiol, glutathione (GSH) (), which yields Cr(V) species that can cause oxidative damage via direct or Fenton-like reactions (; ). Findings in cultured cells commonly guide the design and interpretation of expensive animal studies and are used to determine the mode of action for regulatory purposes. Thus, it is critical to ensure that in vitro models adequately recapitulate the main metabolic processes for Cr(VI) in tissues.
The concept of oxidation state (oxidation number) is then introduced and applied to elements is ion formulae and compound formulae (ionic or covalent).
The two (2) half-reactions are:Notice that both half-reactions are shown as reductions -- the species gains electrons, and is changed to a new form.
Explosions, like the one shown in Daylight and many other movies, are nothing more than oxidation-reduction reactions which happen very rapidly.
10.1.4: Identify whether an element is oxidised or reduced in simple redox reactions, using oxidation numbers. Appropriate reactions to illustrate this can be found in topics 3 and 11. Possible examples include: iron(II) and (III), manganese(II) and (VII), chromium(III) and (VI), copper(I) and (II), oxides of sulphur and oxyacids, halogens and halide ions.
10.1.3: State and explain the relationship between oxidation numbers and the names of compounds. Oxidation numbers in names of compounds are represented by Roman numerals, eg iron (II) oxide, iron (III) oxide.
6. Energy is released in the form of heat and light in combustion, familiarly known as fire, which is the next step up in oxidation reactions in terms of speed and energy output.
See a summary of what "oxidation" can mean in
Scientists have devised a way to unify all three meanings in one, defined as "the increase in oxidation number." This also explains the origin of the word reduction, as it then means "reducing the oxidation number." For a detailed explanation of oxidation numbers see .
You should be able to work out the oxidation state of an element in a compound or ion from the formula, write half-equations identifying the oxidation and reduction processes in redox reactions.
Hydrogen, for example always has an oxidation number of -1 when bonded to a metal (more electropositive element) and +1 when bonded to a more electronegative element (non-metal). Oxygen is always -2 (except when in the form of the peroxide ion when it has an O-O bond giving it an oxidation number of -1). Group 1 and 2 metals usually have an oxidation number of 1+ and 2+ respectively.
the equations are not meant to be balanced, but just focus on the particular oxidation or reduction change and later the full balanced redox equations are given.
There are some elements that virtually always have the same oxidation number and these can be used to calculate the oxidation numbers of the atoms in question.