In chemistry, the oxidation state, or oxidation number, is the hypothetical charge of an atom if 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. Beside nearly-pure ionic bonding, many covalent bonds exhibit a strong ionicity, making oxidation state a useful predictor of charge.
The oxidation state of an atom does not represent the "real" charge on that atom, or any other actual atomic property. This is particularly true of high oxidation states, where the ionization energy required to produce a multiply positive ion is far greater than the energies available in chemical reactions. Additionally, the oxidation states of atoms in a given compound may vary depending on the choice of electronegativity scale used in their calculation. Thus, the oxidation state of an atom in a compound is purely a formalism. It is nevertheless important in understanding the nomenclature conventions of inorganic compounds. Also, several observations regarding chemical reactions may be explained at a basic level in terms of oxidation states.
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 reported to be +9, displayed by iridium in the tetroxoiridium(IX) cation (). It is predicted that even a +10 oxidation state may be achieved by 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. Fe<sup>III</sup><sub>2</sub>O<sub>3</sub>. 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 the oxidation state of an atom, through a chemical reaction, is known as oxidation; a decrease in oxidation state is known as a reduction. Such reactions involve the formal transfer of electrons: a net gain in electrons being a reduction, and a net loss of electrons being oxidation. For pure elements, the oxidation state is zero.
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 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.
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. 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:
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Algorithm of summing bond orders
This algorithm works on Lewis structures and bond graphs of extended (non-molecular) solids:
Applied to a Lewis structure
An example of a Lewis structure with no formal charge,
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illustrates that, in this algorithm, homonuclear bonds are simply ignored (the bond orders are in blue).
Carbon monoxide exemplifies a Lewis structure with formal charges:
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To obtain the oxidation states, the formal charges are summed with the bond-order value taken positively at the carbon and negatively at the oxygen.
Applied to molecular ions, this algorithm considers the actual location of the formal (ionic) charge, as drawn in the Lewis structure. As an example, summing bond orders in the ammonium cation yields −4 at the nitrogen of formal charge +1, with the two numbers adding to the oxidation state of −3:
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The sum of oxidation states in the ion equals its charge (as it equals zero for a neutral molecule).
Also in anions, the formal (ionic) charges have to be considered when nonzero. For sulfate this is exemplified with the skeletal or Lewis structures (top), compared with the bond-order formula of all oxygens equivalent and fulfilling the octet and 8 − N rules (bottom):
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Applied to bond graph
A bond graph in solid-state chemistry is a chemical formula of an extended structure, in which direct bonding connectivities are shown. An example is the perovskite, the unit cell of which is drawn on the left and the bond graph (with added numerical values) on the right:
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We see that the oxygen atom bonds to the six nearest rubidium cations, each of which has 4 bonds to the auride anion. The bond graph summarizes these connectivities. The bond orders (also called bond valences) sum up to oxidation states according to the attached sign of the bond's ionic approximation (there are no formal charges in bond graphs).
Determination of oxidation states from a bond graph can be illustrated on ilmenite, . We may ask whether the mineral contains and , or and . 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. Experimental data show that three metal-oxygen bonds in the octahedron are short and three are long (the metals are off-center). The bond orders (valences), obtained from the bond lengths by the bond valence method, sum up to 2.01 at Fe and 3.99 at Ti; which can be rounded off to oxidation states +2 and +4, respectively:
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Balancing redox
Oxidation states can be useful for balancing chemical equations for oxidation-reduction (or redox) reactions, because the changes in the oxidized atoms have to be balanced by the changes in the reduced atoms. For example, in the reaction of acetaldehyde with Tollens' reagent to form acetic acid (shown below), the carbonyl carbon atom changes its oxidation state from +1 to +3 (loses two electrons). This oxidation is balanced by reducing two cations to (gaining two electrons in total).
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An inorganic example is the Bettendorf reaction using tin dichloride () to prove the presence of arsenite ions in a concentrated HCl extract. When arsenic(III) is present, a brown coloration appears forming a dark precipitate of arsenic, according to the following simplified reaction:
:
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. The simple one-line balancing goes as follows: the two redox couples are written down as they react;
:
One tin is oxidized from oxidation state +2 to +4, a two-electron step, hence 2 is written in front of the two arsenic partners. One arsenic is reduced from +3 to 0, a three-electron step, hence 3 goes in front of the two tin partners. An alternative three-line procedure is to write separately the half-reactions for oxidation and reduction, each balanced with electrons, and then to sum them up such that the electrons cross out. In general, these redox balances (the one-line balance or each half-reaction) need to be checked for the ionic and electron charge sums on both sides of the equation being indeed equal. If they are not equal, suitable ions are added to balance the charges and the non-redox elemental balance.
Appearances
Nominal oxidation states
A nominal oxidation state is a general term with two different definitions:
- Electrochemical oxidation state magnetic or structural data.
- when the bond order has to be ascertained along with an isolated tandem of a heteronuclear and a homonuclear bond. An example is thiosulfate having two possible oxidation states (bond orders are in blue and formal charges in green):
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:The S–S distance measurement in thiosulfate is needed to reveal that this bond order is very close to 1, as in the formula on the left.
Ambiguous/arbitrary oxidation states
- when the electronegativity difference between two bonded atoms is very small (as in ). Two almost equivalent pairs of oxidation states, arbitrarily chosen, are obtained for these atoms.
- when an electronegative p-block atom forms solely homonuclear bonds, the number of which differs from the number of two-electron bonds suggested by rules. Examples are homonuclear finite chains like (the central nitrogen connects two atoms with four two-electron bonds while only three two-electron bonds are required by the 8 − N rule
Fractional oxidation states
Fractional oxidation states are often used to represent the average oxidation state of several atoms of the same element in a structure. For example, the formula of magnetite is , implying an average oxidation state for iron of +. However, this average value may not be representative if the atoms are not equivalent. In a crystal below , two-thirds of the cations are and one-third are , and the formula may be more clearly represented as FeO·.
Likewise, propane, , has been described as having a carbon oxidation state of −. Again, this is an average value since the structure of the molecule is , with the first and third carbon atoms each having an oxidation state of −3 and the central one −2.
An example with true fractional oxidation states for equivalent atoms is potassium superoxide, . The diatomic superoxide ion has an overall charge of −1, so each of its two equivalent oxygen atoms is assigned an oxidation state of −. This ion can be described as a resonance hybrid of two Lewis structures, where each oxygen has an oxidation state of 0 in one structure and −1 in the other.
For the cyclopentadienyl anion , the oxidation state of C is −1 + − = −. The −1 occurs because each carbon is bonded to one hydrogen atom (a less electronegative element), and the − because the total ionic charge of −1 is divided among five equivalent carbons. Again this can be described as a resonance hybrid of five equivalent structures, each having four carbons with oxidation state −1 and one with −2.
:{| class="wikitable"
|+ Examples of fractional oxidation states for carbon
|-
! Oxidation state !! Example species
|-
| − ||
|-
| − ||
|-
| + ||
|}
Finally, fractional oxidation numbers are not used in the chemical nomenclature. For example the red lead is represented as lead(II,IV) oxide, showing the oxidation states of the two nonequivalent lead atoms.
Elements with multiple oxidation states
Most elements have more than one possible oxidation state. For example, carbon has nine possible integer oxidation states from −4 to +4:
:{| class="wikitable"
|+ Integer oxidation states of carbon
|-
! Oxidation state !! Example compound
|-
| −4 ||
|-
| −3 ||
|-
| −2 || ,
|-
| −1 || , ,
|-
| 0 || ,
|-
| +1 || ,
|-
| +2 || ,
|-
| +3 || ,
|-
| +4 || ,
|}
Oxidation state in metals
Many compounds with luster and electrical conductivity maintain a simple stoichiometric formula, such as the golden TiO, blue-black or coppery , all of obvious oxidation state. Ultimately, assigning the free metallic electrons to one of the bonded atoms is not comprehensive and can yield unusual oxidation states. Examples are the LiPb and ordered alloys, the composition and structure of which are largely determined by atomic size and packing factors. Should oxidation state be needed for redox balancing, it is best set to 0 for all atoms of such an alloy.
List of oxidation states of the elements
This is a list of known oxidation states of the chemical elements, excluding nonintegral values. The most common states appear in bold. The table is based on that of Greenwood and Earnshaw, with additions noted. Every element exists in oxidation state 0 when it is the pure non-ionized element in any phase, whether monatomic or polyatomic allotrope. The column for oxidation state 0 only shows elements known to exist in oxidation state 0 in compounds.
Early forms (octet rule)
A figure with a similar format was used by Irving Langmuir in 1919 in one of the early papers about the octet rule. The periodicity of the oxidation states was one of the pieces of evidence that led Langmuir to adopt the rule.
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Use in nomenclature
The oxidation state in compound naming for transition metals and lanthanides and actinides is placed either as a right superscript to the element symbol in a chemical formula, such as Fe<sup>III</sup> or in parentheses after the name of the element in chemical names, such as iron(III). For example, is named iron(III) sulfate and its formula can be shown as . This is because a sulfate ion has a charge of −2, so each iron atom takes a charge of +3.
History of the oxidation state concept
Early days
Oxidation itself was first studied by Antoine Lavoisier, who defined it as the result of reactions with oxygen (hence the name). The term has since been generalized to imply a formal loss of electrons. Oxidation states, called oxidation grades by Friedrich Wöhler in 1835, were one of the intellectual stepping stones that Dmitri Mendeleev used to derive the periodic table. William B. Jensen gives an overview of the history up to 1938.
Use in nomenclature
When it was realized that some metals form two different binary compounds with the same nonmetal, the two compounds were often distinguished by using the ending -ic for the higher metal oxidation state and the ending -ous for the lower. For example, is ferric chloride and is ferrous chloride. This system is not very satisfactory (although sometimes still used) because different metals have different oxidation states which have to be learned: ferric and ferrous are +3 and +2 respectively, but cupric and cuprous are +2 and +1, and stannic and stannous are +4 and +2. Also, there was no allowance for metals with more than two oxidation states, such as vanadium with oxidation states +2, +3, +4, and +5. and adopted by IUPAC in 1940. Thus, was written as iron(II) chloride rather than ferrous chloride. The Roman numeral II at the central atom came to be called the "Stock number" (now an obsolete term), and its value was obtained as a charge at the central atom after removing its ligands along with the electron pairs they shared with it. He used it for the value (synonymous with the German term Wertigkeit) previously termed "valence", "polar valence" or "polar number" in English, or "oxidation stage" or indeed the "state of oxidation". Since 1938, the term "oxidation state" has been connected with electrochemical potentials and electrons exchanged in redox couples participating in redox reactions. By 1948, IUPAC used the 1940 nomenclature rules with the term "oxidation state", instead of the original A full acceptance of this suggestion was complicated by the fact that the Pauling electronegativities as such depend on the oxidation state and that they may lead to unusual values of oxidation states for some transition metals. In 1990 IUPAC resorted to a postulatory (rule-based) method to determine the oxidation state. This was complemented by the synonymous term oxidation number as a descendant of the Stock number introduced in 1940 into the nomenclature. However, the terminology using "ligands"