Oxidation state

In some cases, the average oxidation state of an element is a fraction, such as 8⁄3 for iron in magnetite . The highest known oxidation state is reported to be +9 in the tetroxoiridium(IX) cation . It is predicted that even a +10 oxidation state may be achievable by platinum in the tetroxoplatinum(X) cation . The lowest oxidation state is −4, as for carbon in methane or for chromium in [Cr(CO) 4 ] 4−.

The most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and transition metal carbonyl complexes.

This approach yields correct oxidation states in oxides and hydroxides of any single element, and in acids such as H 2 SO 4 or H 2 Cr 2 O 7 . Its coverage can be extended either by a list of exceptions or by assigning priority to the postulates.

Chromic acid features chromium in an oxidation state of +6 (or VI).

In some cases, the average oxidation state of an element is a fraction, such as 8⁄3 for iron in magnetite . The highest known oxidation state is reported to be +9 in the tetroxoiridium(IX) cation . It is predicted that even a +10 oxidation state may be achievable by platinum in the tetroxoplatinum(X) cation . The lowest oxidation state is −4, as for carbon in methane or for chromium in [Cr(CO) 4 ] 4−.

The most common oxidation states of platinum are +2 and +4. The +1 and +3 oxidation states are less common, and are often stabilized by metal bonding in bimetallic (or polymetallic) species.

It also covers iodides, sulfides and similar simple salts of these metals.

An iodide ion is the ion I − . Compounds with iodine in formal oxidation state −1 are called iodides. This page is for the iodide ion and its salts, not organoiodine compounds.

It covers all oxoacids of any central atom (and all their fluoro-, chloro- and bromo-relatives), as well as salts of such acids with group 1 and 2 metals.

Inorganic oxyacids typically have a chemical formula of type H m XO n, where X is an atom functioning as a central atom, whereas parameters m and n depend on the oxidation state of the element X. In most cases, the element X is a nonmetal, but some metals, for example chromium and manganese, can form oxyacids when occurring at their highest oxidation states.

4) Group 1 and group 2 metals in compounds have OS = +1 and +2, respectively.

that is, this orbital contains its full complement of two electrons, which these elements readily lose to form cations with charge +2, and an oxidation state of +2.

In some cases, the average oxidation state of an element is a fraction, such as 8⁄3 for iron in magnetite . The highest known oxidation state is reported to be +9 in the tetroxoiridium(IX) cation . It is predicted that even a +10 oxidation state may be achievable by platinum in the tetroxoplatinum(X) cation . The lowest oxidation state is −4, as for carbon in methane or for chromium in [Cr(CO) 4 ] 4−.

Chromium exhibits a wide range of oxidation states, but chromium being ionized into a cation with a positive 3 charge serves as chromium's most stable ionic state.

Much later, it was realized that the substance, upon being oxidized, loses electrons, and the meaning was extended to include other reactions in which electrons are lost, regardless of whether oxygen was involved.

Oxidation is better defined as an increase in oxidation state, and reduction as a decrease in oxidation state.

In some cases, the average oxidation state of an element is a fraction, such as 8⁄3 for iron in magnetite . The highest known oxidation state is reported to be +9 in the tetroxoiridium(IX) cation . It is predicted that even a +10 oxidation state may be achievable by platinum in the tetroxoplatinum(X) cation . The lowest oxidation state is −4, as for carbon in methane or for chromium in [Cr(CO) 4 ] 4−.

Iridium tetroxide (IrO 4, Iridium(VIII) oxide) is a binary compound of oxygen and iridium in oxidation state +VIII.

Analogously for transition-metal compounds; CrO(O 2 ) 2 on the left has a total of 36 valence electrons (18 pairs to be distributed), and Cr(CO) 6 on the right has 66 valence electrons (33 pairs): The algorithm contains a caveat, which concerns rare cases of transition-metal complexes with a type of ligand that is reversibly bonded as a Lewis acid (as an acceptor of the electron pair from the transition metal); termed a "Z-type" ligand in Green’s covalent bond classification method.

In the oxidation state +2 the ions have the electronic configuration [ ] d 10.

One early example is the O 2 S−RhCl(CO)(PPh 3 ) 2 complex with SO 2 as the reversibly-bonded acceptor ligand (released upon heating).

In terms of electron-counting formalism, the sulfur atom has an oxidation state of +4 and a formal charge of +1.

3) Fluorine in compounds has OS = −1; this extends to chlorine and bromine only when not bonded to a lighter halogen, oxygen or nitrogen.

The second example also involves a reduction in oxidation state, which can also be achieved by reducing a higher chloride using hydrogen or a metal as a reducing agent.

3) Fluorine in compounds has OS = −1; this extends to chlorine and bromine only when not bonded to a lighter halogen, oxygen or nitrogen.

It combines with metals, nonmetals, metalloids, and most noble gases, and usually assumes an oxidation state of −1.

Its crystal structure has each metal atom bonded to six oxygens and each of the equivalent oxygens to two irons and two titaniums, as in the bond graph below.

Like the other group 8 elements, ruthenium and osmium, iron exists in a wide range of oxidation states, −2 to +7, although +2 and +3 are the most common.

Analogously for transition-metal compounds; CrO(O 2 ) 2 on the left has a total of 36 valence electrons (18 pairs to be distributed), and Cr(CO) 6 on the right has 66 valence electrons (33 pairs):

Cr(CO) 6 is zerovalent, meaning that Cr has an oxidation state of zero, and it is a homoleptic complex, which means that all the ligands are identical.

In inorganic nomenclature, the oxidation state is represented by a Roman numeral placed after the element name inside a parenthesis or as a superscript after the element symbol.

They are also used in the IUPAC nomenclature of inorganic chemistry, for the oxidation number of cations which can take on several different positive charges.

The algorithm contains a caveat, which concerns rare cases of transition-metal complexes with a type of ligand that is reversibly bonded as a Lewis acid (as an acceptor of the electron pair from the transition metal); termed a "Z-type" ligand in Green’s covalent bond classification method.

It was published by M. L. H. Green in the mid-1990s as a solution for the need to describe covalent compounds such as organometallic complexes in a way that is not prone to limitations resulting from the definition of oxidation state.

The oxidation state of gold in its compounds ranges from −1 to +5, but Au(I) and Au(III) dominate its chemistry.

This algorithm is performed on a Lewis structure (a formula that shows all valence electrons).

In effect, there are possibly seven valence electrons (4s 2 3d 5 ) outside the argon-like core; this is consistent with the chemical fact that manganese can have an oxidation state as high as +7 (in the permanganate ion: ).

*Electrochemical oxidation state; it represents a molecule or ion in the Latimer diagram or Frost diagram for its redox-active element.

A Frost diagram or Frost–Ebsworth diagram is a type of graph used by inorganic chemists in electrochemistry to illustrate the relative stability of a number of different oxidation states of a particular substance.

Its summary formula, HNO 3, corresponds to two structural isomers; the peroxynitrous acid in the above figure and the more stable nitric acid.

Reaction with non-metallic elements, with the exceptions of nitrogen, oxygen, noble gases, silicon, and halogens other than iodine, usually oxidizes them to their highest oxidation states as acids with the formation of nitrogen dioxide for concentrated acid and nitric oxide for dilute acid.

Additionally, oxidation states of atoms in a given compound may vary depending on the choice of electronegativity scale used in their calculation.

While this approach has the advantage of simplicity, it is clear that the electronegativity of an element is not an invariable atomic property and, in particular, increases with the oxidation state of the element.

Here three tin atoms are oxidized from oxidation state +2 to +4, yielding six electrons that reduce two arsenic atoms from oxidation state +3 to 0.

It is obtained chiefly from the mineral cassiterite, which contains stannic oxide, SnO 2 . Tin shows a chemical similarity to both of its neighbors in group 14, germanium and lead, and has two main oxidation states, +2 and the slightly more stable +4. Tin is the 49th most abundant element and has, with 10 stable isotopes, the largest number of stable isotopes in the periodic table, thanks to its magic number of protons.

A key step is drawing the Lewis structure of the molecule (neutral, cationic, anionic): atom symbols are arranged so that pairs of atoms can be joined by single two-electron bonds as in the molecule (a sort of "skeletal" structure), and the remaining valence electrons are distributed such that sp atoms obtain an octet (duet for hydrogen) with priority that increases with electronegativity.

In 1904 Richard Abegg was one of the first to extend the concept of coordination number to a concept of valence in which he distinguished atoms as electron donors or acceptors, leading to positive and negative valence states that greatly resemble the modern concept of oxidation states.

It also covers iodides, sulfides and similar simple salts of these metals.

Molybdenum disulfide (MoS 2 ) consists of separated sulfide centers, in association with molybdenum in the formal +4 oxidation state (that is, Mo 4+ and two S 2− ).