Oxidation state

The oxidation state, or oxidation number, is the hypothetical charge of an atom if all of its bonds to different atoms were 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. While fully ionic bonds are not found in nature, many bonds exhibit strong ionicity, making oxidation state a useful predictor of charge.

Number that describes the degree of oxidation of an atom in a chemical compound; the hypothetical charge that an atom would have if all bonds to atoms of different elements were fully ionic

The oxidation state of an atom does not represent the “real” formal 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

8/3 for iron in magnetiteFe
3
O
4
(see below). The highest known oxidation state is reported to be +9 in the tetroxoiridium(IX) cation (IrO+
4
).[1] It is predicted that even a +12 oxidation state may be achievable by uranium in the unusual hexoxide UO6.[2] The lowest oxidation state is −5, as for boron in Al3BC.[3]

In inorganic nomenclature, the oxidation state is represented by a Roman numeral placed after the element name inside the parenthesis or as a superscript after the element symbol, e.g. Iron(III) oxide.

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.

. . . Oxidation state . . .

IUPAC has published a “Comprehensive definition of the term oxidation state (IUPAC Recommendations 2016)”.[4] It is a distillation of an IUPAC technical report “Toward a comprehensive definition of oxidation state” from 2014.[5] The current IUPAC Gold Book definition of oxidation state is:

Oxidation state of an atom is the charge of this atom after ionic approximation of its heteronuclear bonds…

IUPAC[6]

and the term oxidation number is nearly synonymous.[7]

The underlying principle is that the ionic charge is “the oxidation state of an atom, after ionic approximation of its bonds”,[8] where ionic approximation means, hypothesizing that all bonds are ionic. Several criteria were considered for the ionic approximation:

  1. Extrapolation of the bond’s polarity;
    1. from the electronegativity difference,
    2. from the dipole moment, and
    3. from quantum‐chemical calculations of charges.
  2. Assignment of electrons according to the atom’s contribution to the bonding MO[8][9]/ the electron’s allegiance in a LCAO–MO model.[10]

In a bond between two different elements, the bond’s electrons are assigned to its main atomic contributor/higher electronegativity; in a bond between two atoms of the same element, the electrons are divided equally. This is because 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 PtH4−2, for Pauling and Mulliken scales.[11] 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,[8] 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:

Electronegativity using the Allen scale
Group  1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
 Period
1 H
2.300
He
4.160
2 Li
0.912
Be
1.576
B
2.051
C
2.544
N
3.066
O
3.610
F
4.193
Ne
4.787
3 Na
0.869
Mg
1.293
Al
1.613
Si
1.916
P
2.253
S
2.589
Cl
2.869
Ar
3.242
4 K
0.734
Ca
1.034
Sc
1.19
Ti
1.38
V
1.53
Cr
1.65
Mn
1.75
Fe
1.80
Co
1.84
Ni
1.88
Cu
1.85
Zn
1.588
Ga
1.756
Ge
1.994
As
2.211
Se
2.424
Br
2.685
Kr
2.966
5 Rb
0.706
Sr
0.963
Y
1.12
Zr
1.32
Nb
1.41
Mo
1.47
Tc
1.51
Ru
1.54
Rh
1.56
Pd
1.58
Ag
1.87
Cd
1.521
In
1.656
Sn
1.824
Sb
1.984
Te
2.158
I
2.359
Xe
2.582
6 Cs
0.659
Ba
0.881
Lu
1.09
Hf
1.16
Ta
1.34
W
1.47
Re
1.60
Os
1.65
Ir
1.68
Pt
1.72
Au
1.92
Hg
1.765
Tl
1.789
Pb
1.854
Bi
2.01
Po
2.19
At
2.39
Rn
2.60
7 Fr
0.67
Ra
0.89
See also: Electronegativities of the elements (data page)

. . . Oxidation state . . .

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. . . Oxidation state . . .

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