In chemistry, the oxidation state, or oxidation number, is the hypothetical charge of an atom assuming all of its bonds to other atoms are fully ionic. It describes the degree of oxidation (loss of electrons) of an atom in a chemical compound. Conceptually, the oxidation state may be positive, negative or zero. In a binary compound like , an oxidation state of +1 for K implies it has donated an electron to Cl, which will have an oxidation state of -1. Oxidation state is useful for estimating the distribution of electric charge within molecules, especially for those containing primarily ionic bonds, and less so for those with primarily covalent bonds.
The oxidation state of an atom is capable of qualitative explanation of patterns in chemical reactions. It does not represent an atom's actual charge. The oxidation states of atoms in a given compound may vary depending on the electronegativity scale used. It is also relevant to the nomenclature conventions of inorganic compounds.
Oxidation states are typically represented by integers which may be positive, zero, or negative. In some cases, the average oxidation state of an element is a fraction, such as for iron in magnetite (see below). The highest known oxidation state is +9, displayed by iridium in the tetroxoiridium(IX) cation (). An even higher +10 is predicted to be possible for platinum in tetroxoplatinum(X), . The lowest oxidation state is −5, as for boron in and gallium in pentamagnesium digallide ().
In Stock nomenclature, which is commonly used for inorganic compounds, the oxidation state is represented by a Roman numeral placed after the element name inside parentheses, e.g. Iron(III) oxide, or as a superscript after the element symbol, e.g. FeIII2O3. The term oxidation was first used by Antoine Lavoisier to signify the reaction of a substance with oxygen. 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. The increase in an atom's oxidation state through a chemical reaction is referred to as oxidation, and a decrease in oxidation state reduction; the two always occur as pairs in redox reactions. For pure elements, the oxidation state is zero.
Oxidation state, which idealizes bonds as being wholly ionic, can be understood as an opposite to formal charge, which assumes bonds are wholly covalent. Both can be useful for predicting the charge distributions of molecules.
Overview Oxidation numbers are assigned to elements in a molecule such that the overall sum is zero in a neutral molecule. The number indicates the degree of oxidation of each element caused by molecular bonding. In ionic compounds, the oxidation numbers are the same as the element's ionic charge. Thus for KCl, potassium is assigned +1 and chlorine is assigned −1. The complete set of rules for assigning oxidation numbers are discussed in the following sections.
Oxidation numbers are fundamental to the chemical nomenclature of ionic compounds. For example, Cu compounds with Cu oxidation state +2 are called cupric and those with state +1 are cuprous.Siebring, B. R., Schaff, M. E. (1980). General Chemistry. United States: Wadsworth Publishing Company. The oxidation numbers of elements allow predictions of chemical formula and reactions, especially oxidation-reduction reactions. The oxidation numbers of the most stable chemical compounds follow trends in the periodic table.Gray, H. B., Haight, G. P. (1967). Basic Principles of Chemistry. Netherlands: W. A. Benjamin.
IUPAC definition
International Union of Pure and Applied Chemistry (IUPAC) has published a "Comprehensive definition of oxidation state (IUPAC Recommendations 2016)". It is a distillation of an IUPAC technical report: "Toward a comprehensive definition of oxidation state". According to the IUPAC Gold Book: "The oxidation state of an atom is the charge of this atom after ionic approximation of its heteronuclear bonds." The term oxidation number is nearly synonymous.
The ionic approximation means extrapolating bonds to ionic. Several criteria were considered for the ionic approximation:
Extrapolation of the bond's polarity; Assignment of electrons according to the atom's contribution to the bonding Molecular orbital (MO) or the electron's allegiance in a LCAO–MO model.
In a bond between two different elements, the bond's electrons are assigned to its main atomic contributor typically of higher electronegativity; in a bond between two atoms of the same element, the electrons are divided equally. Most electronegativity scales depend on the atom's bonding state, which makes the assignment of the oxidation state a somewhat circular argument. For example, some scales may turn out unusual oxidation states, such as −6 for platinum in , for Pauling and Mulliken scales. The dipole moments would sometimes also turn out abnormal oxidation numbers, such as in CO and NO, which are oriented with their positive end towards oxygen. Therefore, this leaves the atom's contribution to the bonding MO, the atomic-orbital energy, and from quantum-chemical calculations of charges, as the only viable criteria with cogent values for ionic approximation. However, for a simple estimate for the ionic approximation, we can use Allen electronegativities, as only that electronegativity scale is truly independent of the oxidation state, as it relates to the average valence‐electron energy of the free atom:
Determination While introductory levels of chemistry teaching use postulated oxidation states, the IUPAC recommendation and the Gold Book entry list two entirely general algorithms for the calculation of the oxidation states of elements in chemical compounds.
Simple approach without bonding considerations Introductory chemistry uses postulates: the oxidation state for an element in a chemical formula is calculated from the overall charge and postulated oxidation states for all the other atoms.
A simple example is based on two postulates,
OS = +1 for hydrogen OS = −2 for oxygen
where OS stands for oxidation state. This approach yields correct oxidation states in oxides and hydroxides of any single element, and in acids such as sulfuric acid () or dichromic acid (). Its coverage can be extended either by a list of exceptions or by assigning priority to the postulates. The latter works for hydrogen peroxide () where the priority of rule 1 leaves both oxygens with oxidation state −1.
Additional postulates and their ranking may expand the range of compounds to fit a textbook's scope. As an example, one postulatory algorithm from many possible; in a sequence of decreasing priority:
An element in a free form has OS = 0. In a compound or ion, the sum of the oxidation states equals the total charge of the compound or ion. Fluorine in compounds has OS = −1; this extends to chlorine and bromine only when not bonded to a lighter halogen, oxygen or nitrogen. Group 1 and group 2 metals in compounds have OS = +1 and +2, respectively. Hydrogen has OS = +1 but adopts −1 when bonded as a hydride to metals or metalloids. Oxygen in compounds has OS = −2 but only when not bonded to oxygen (e.g. in peroxides) or fluorine.
This set of postulates covers oxidation states of fluorides, chlorides, bromides, oxides, hydroxides, and hydrides of any single element. 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. It also covers iodides, sulfides, and similar simple salts of these metals.
Algorithm of assigning bonds This algorithm is performed on a Lewis structure (a diagram that shows all valence electrons). Oxidation state equals the charge of an atom after each of its heteronuclear bonds has been assigned to the more electronegative partner of the bond (except when that partner is a reversibly bonded Lewis-acid ligand) and homonuclear bonds have been divided equally:
where each "—" represents an electron pair (either shared between two atoms or solely on one atom), and "OS" is the oxidation state as a numerical variable.
After the electrons have been assigned according to the vertical red lines on the formula, the total number of valence electrons that now "belong" to each atom is subtracted from the number of valence electrons of the neutral atom (such as 5 for nitrogen in group 15) to yield that atom's oxidation state.
This example shows the importance of describing the bonding. Its summary formula, , corresponds to two structural isomers; the peroxynitrous acid in the above figure and the more stable nitric acid. With the formula , the simple approach without bonding considerations yields −2 for all three oxygens and +5 for nitrogen, which is correct for nitric acid. For the peroxynitrous acid, however, both oxygens in the O–O bond have OS = −1, and the nitrogen has OS = +3, which requires a structure to understand.
Organic compounds are treated in a similar manner; exemplified here on functional groups occurring in between methane () and carbon dioxide ():
Analogously for transition-metal compounds; on the left has a total of 36 valence electrons (18 pairs to be distributed), and hexacarbonylchromium () on the right has 66 valence electrons (33 pairs):
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 a priority that increases in proportion with electronegativity. In some cases, this leads to alternative formulae that differ in bond orders (the full set of which is called the resonance formulas). Consider the sulfate anion () with 32 valence electrons; 24 from oxygens, 6 from sulfur, 2 of the anion charge obtained from the implied cation. The bond orders to the terminal oxygens do not affect the oxidation state so long as the oxygens have octets. Already the skeletal structure, top left, yields the correct oxidation states, as does the Lewis structure, top right (one of the resonance formulas):
The bond-order formula at the bottom is closest to the reality of four equivalent oxygens each having a total bond order of 2. That total includes the bond of order to the implied cation and follows the 8 − N rule requiring that the main-group atom's bond-order total equals 8 − N valence electrons of the neutral atom, enforced with a priority that proportionately increases with electronegativity.
This algorithm works equally for molecular cations composed of several atoms. An example is the ammonium cation of 8 valence electrons (5 from nitrogen, 4 from hydrogens, minus 1 electron for the cation's positive charge):
Drawing Lewis structures with electron pairs as dashes emphasizes the essential equivalence of bond pairs and lone pairs when counting electrons and moving bonds onto atoms. Structures drawn with electron dot pairs are of course identical in every way:
The algorithm's caveat 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. The caveat originates from the simplifying use of electronegativity instead of the MO-based electron allegiance to decide the ionic sign. One early example is the complex with sulfur dioxide () as the reversibly-bonded acceptor ligand (released upon heating). The Rh−S bond is therefore extrapolated ionic against Allen electronegativities of rhodium and sulfur, yielding oxidation state +1 for rhodium:
Algorithm of summing bond orders