Oxidation state

The oxidation state, sometimes referred to as oxidation number, describes the degree of oxidation (loss of electrons) of an atom in a chemical compound. Conceptually, the oxidation state, which may be positive, negative or zero, is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic, with no covalent component. This is never exactly true for real bonds.

The term oxidation was first used by Antoine Lavoisier to signify 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.

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 magnetite (Fe
3
O
4
). The highest known oxidation state is reported to be +9 in the tetroxoiridium(IX) cation (IrO+
4
).[1] It is predicted that even a +10 oxidation state may be achievable by platinum in the tetroxoplatinum(X) cation (PtO2+
4
).[2] The lowest oxidation state is −4, as for carbon in methane or for chromium[3] in [Cr(CO)4]4−.

The increase in oxidation state of an atom, through a chemical reaction, is known as an 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 an oxidation. For pure elements, the oxidation state is zero.

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, 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, a number of observations pertaining to chemical reactions may be explained at a basic level in terms of oxidation states.

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.

IUPAC definition of oxidation state

Oxidation state

A "Comprehensive definition of the term oxidation state (IUPAC Recommendations 2016)" has been published with free access.[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[6] is:

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

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

The underlying principle is that the ionic signs for two atoms that are bonded are deduced from the electron distribution in a LCAO–MO model. In a bond between two different elements, the bond's electrons are assigned to its main atomic contributor; in a bond between two atoms of the same element, the electrons are divided equally. In practical use, the sign of the ionic approximation follows Allen electronegativities:

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
1.59
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
1.52
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
1.76
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)

Determination of oxidation state

While introductory levels of chemistry teaching use postulated oxidation states, the IUPAC recommendation[4] and the Gold Book entry[6] 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,

  1. OS = +1 for hydrogen
  2. 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 H2SO4 or H2Cr2O7. Its coverage can be extended either by a list of exceptions or by assigning priority to the postulates. The latter works for H2O2 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:

  1. An element in a free form has OS = 0.
  2. In a compound or ion, the oxidation states' sum equals the total charge of the compound or ion.
  3. Fluorine in compounds has OS = −1; this extends to chlorine and bromine only when not bonded to a lighter halogen, oxygen or nitrogen.
  4. Group 1 and group 2 metals in compounds have OS = +1 and +2, respectively.
  5. Hydrogen has OS = +1, but adopts −1 when bonded as a hydride to metals or metalloids.
  6. Oxygen in compounds has OS = −2.

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 formula that shows all valence electrons). Oxidation state equals the charge of an atom after its heteronuclear bonds have been assigned to the more electronegative partner (except when that partner is a reversibly bonded Lewis-acid ligand) and homonuclear bonds have been divided equally:

1oxstate
1oxstate

where "—" is an electron pair, 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 are subtracted from the number N 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, HNO3, corresponds to two structural isomers; the peroxynitrous acid in the above figure and the more stable nitric acid. With the formula HNO3, 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, the two oxygens in the O–O bond each 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 CH4 and CO2:

3oxstate
3oxstate

Analogously for transition-metal compounds; CrO(O2)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):

2oxstate
2oxstate

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 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 (SO2−
4
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 have no effect on 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):

7oxstate
7oxstate

The bond-order formula at 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 1/2 to the implied cation and follows the 8 − N rule[5] requiring that the main-group atom’s bond order equals 8 minus N valence electrons of the neutral atom, enforced with priority that 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):

5oxstate
5oxstate

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:

4oxstate
4oxstate

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.[4] One early example is the O2S−RhCl(CO)(PPh3)2 complex[8] with SO2 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:

8oxstate
8oxstate

Algorithm of summing bond orders

This algorithm works on Lewis structures and on bond graphs of extended (non-molecular) solids:

Oxidation state is obtained by summing the heteronuclear-bond orders at the atom as positive if that atom is the electropositive partner in a particular bond and as negative if not, and the atom’s formal charge (if any) is added to that sum.

Applied to a Lewis structure

An example of a Lewis structure with no formal charge,

9oxstate
9oxstate

illustrates that, in this algorithm, homonuclear bonds are simply ignored (notice the bond orders in blue).

Carbon monoxide exemplifies a Lewis structure with formal charges:

10oxstate
10oxstate

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:

11oxstate
11oxstate

Notice that 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):

13oxstate
13oxstate

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 AuORb3 perovskite, the unit cell of which is drawn on the left and the bond graph (with added numerical values) on the right:

14oxstate
14oxstate

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, FeTiO3. We may ask whether the mineral contains Fe2+ and Ti4+, or Fe3+ and Ti3+. 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:

15oxstate
15oxstate

Nominal oxidation states

A nominal oxidation state is a general term for two specific purpose-oriented values:

16oxstate
16oxstate
  • Systematic oxidation state; it is chosen from close alternatives for pedagogical reasons of descriptive chemistry. An example is the oxidation state of phosphorus in H3PO3 (which is in fact the diprotic HPO(OH)2) taken nominally as +3, while Allen electronegativities of phosphorus and hydrogen suggest +5 by a narrow margin that makes the two alternatives almost equivalent:
17oxstate
17oxstate

Both alternative oxidation states of phosphorus make chemical sense, depending on the chemical property or reaction we wish to emphasize. In contrast, their average (+4) does not.

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 Ag+ cations to Ag0 (gaining two electrons in total).

Redox eqn 1
Redox eqn 1

An inorganic example is the Bettendorf reaction using SnCl2 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:

2 As3+ + 3 Sn2+ → 2 As0 + 3 Sn4+

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;

As3+ + Sn2+ ⇌ As0 + Sn4+.

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 for 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.

Ambiguous oxidation states

Lewis formulae are fine rule-based approximations of chemical reality, as indeed are Allen electronegativities. Still, oxidation states may seem ambiguous when their determination is not straightforward. Rule-based oxidation states feel ambiguous when only experiment can decide. There are also truly dichotomous values to be decided by mere convenience.

Oxidation-state determination from resonance formulas is not straightforward

Seemingly ambiguous oxidation states are obtained on a set of resonance formulas of equal weights for a molecule of heteronuclear bonds where the atom connectivity does not correspond to the number of two-electron bonds dictated by the 8 − N rule. An example is S2N2 where four resonance formulas featuring one S=N double bond have oxidation states +2 and +4 on the two sulfur atoms, to be averaged to +3 because the two sulfur atoms are equivalent in this square-shaped molecule.

A physical measurement is needed to decide the oxidation state

  • This happens when a non-innocent ligand is present, of hidden or unexpected redox properties that could otherwise be assigned to the central atom. An example is the nickel dithiolate complex, Ni(S
    2
    C
    2
    H
    2
    )2−
    2
    .[5]:1056–1057
  • When the redox ambiguity of a central atom and ligand yields dichotomous oxidation states of close stability, thermally induced tautomerism may result, as exemplified by manganese catecholate, Mn(C6H4O2)3.[5]:1057–1058 Assignment of such oxidation states in general requires spectroscopic,[9] magnetic or structural data.
  • When the bond order has to be ascertained along an isolated tandem of a heteronuclear and a homonuclear bond. An example is S
    2
    O2−
    3
    with two oxidation-state alternatives (note bond orders in blue and formal charges in green):
21oxstate
21oxstate
The S–S distance in thiosulfate is needed to reveal that this bond order is very close to 1, as in the formula on the left.

Truly ambiguous oxidation states occur

  • When the electronegativity difference between two bonded atoms is very small (as in H3PO3 above). Two almost equivalent pairs of oxidation states, open for a choice, 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 N
    3
    (the central nitrogen connects two atoms while three two-electron bonds are required by 8 − N rule) or I
    3
    (the central iodine connects two atoms while one two-electron bond fulfills the 8 − N rule). A sensible approach is to distribute the ionic charge over the two outer atoms.[5] Such a placement of charges in a polysulfide S2−
    n
    (where all inner sulfurs form two bonds, fulfilling the 8 − N rule) follows already from its Lewis structure.[5]
  • When the isolated tandem of a heteronuclear and a homonuclear bond leads to a bonding compromise in between two Lewis structures of limiting bond orders. An example here is N2O:
18oxstate
18oxstate
The typically-used oxidation state of nitrogen in N2O is +1, which also obtains for both nitrogens by a molecular orbital approach.[10] It is worth noting that the formal charges on the right comply with electronegativities, and this implies an added ionic bonding contribution. Indeed, the estimated N−N and N−O bond orders are 2.76 and 1.9, respectively,[5] approaching the formula of integer bond orders that would include the ionic contribution explicitly as a bond (in green):
19oxstate
19oxstate
Conversely, formal charges against electronegativities in a Lewis structure decrease the bond order of the corresponding bond. An example is carbon monoxide with a bond-order estimate of 2.6.[11]

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:

Integer oxidation states of carbon
Oxidation state Example compound
−4 CH
4
−3 C
2
H
6
−2 C
2
H
4
, CH
3
Cl
−1 C
2
H
2
, C
6
H
6
, (CH
2
OH)
2
0 HCHO, CH
2
Cl
2
+1 OCHCHO, CHCl
2
CHCl
2
+2 HCOOH, CHCl
3
+3 HOOCCOOH, C
2
Cl
6
+4 CCl
4
, CO
2

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 Fe
3
O
4
, implying an average oxidation state for iron of +8/3.[12]:81–82 However, this average value may not be representative if the atoms are not equivalent. In a Fe
3
O
4
crystal below 120 K (−153 °C), two-thirds of the cations are Fe3+
and one-third are Fe2+
, and the formula may be more specifically represented as FeO·Fe
2
O
3
.[13]

Likewise, propane, C
3
H
8
, has been described as having a carbon oxidation state of −8/3.[14] Again, this is an average value since the structure of the molecule is H
3
C−CH
2
−CH
3
, 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, KO
2
. The diatomic superoxide ion O
2
has an overall charge of −1, so each of its two equivalent oxygen atoms is assigned an oxidation state of −1/2. This ion can be described as a resonance hybrid of two Lewis structures, where each oxygen has oxidation state 0 in one structure and −1 in the other.

For the cyclopentadienyl anion C
5
H
5
, the oxidation state of C is −1 + −1/5 = −6/5. The −1 occurs because each carbon is bonded to one hydrogen atom (a less electronegative element), and the −1/5 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.

Examples of fractional oxidation states for carbon
Oxidation state Example species
6/5 C
5
H
5
6/7 C
7
H+
7
+3/2 C
4
O2−
4

Use in nomenclature

The oxidation state in compound naming is placed either as a right superscript to the element symbol in a chemical formula, such as FeIII, or in parentheses after the name of the element in chemical names, such as iron(III). For example, Fe
2
(SO
4
)
3
is named iron(III) sulfate and its formula can be shown as FeIII
2
(SO
4
)
3
. This is because a sulfate ion has a charge of −2, so each iron atom takes a charge of +3. Note that fractional oxidation numbers should not be used in naming.[15]:66 Minium, Pb
3
O
4
, is represented as lead(II,IV) oxide, showing the actual two oxidation states of the nonequivalent lead atoms.

Oxidation state in metals

Many compounds with luster and electrical conductivity maintain a simple stoichiometric formula; such as the golden TiO, blue-black RuO2 or coppery ReO3, all of obvious oxidation state. Ultimately, however, the assignment of the free metallic electrons to one of the bonded atoms has its limits and leads to unusual oxidation states. Simple examples are the LiPb and Cu3Au 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.

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).[16][17] The term has since been generalized to imply a formal loss of electrons. Oxidation states, called oxidation grades by Friedrich Wöhler in 1835,[18] were one of the intellectual stepping stones that Dmitri Mendeleev used to derive the periodic table. Jensen[19] 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, FeCl3 is ferric chloride and FeCl2 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.[12]:84

This system has been largely replaced by one suggested by Alfred Stock in 1919[20] and adopted[21] by IUPAC in 1940. Thus, FeCl2 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.[15]:147

Development towards the current concept

The term "oxidation state" in English chemical literature was popularized by Wendell Mitchell Latimer in his 1938 book about electrochemical potentials.[22] He used it for the value (synonymous with the German term Wertigkeit) previously termed "valence", "polar valence" or "polar number"[23] in English, or "oxidation stage" or indeed[24][25] 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",[26][27] instead of the original[21] valency. In 1948 Linus Pauling proposed that oxidation number could be determined by extrapolating bonds to being completely ionic in the direction of electronegativity.[28] 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.[29] 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"[15]:147 gave the impression that oxidation number might be something specific to coordination complexes. This situation and the lack of a real single definition generated numerous debates about the meaning of oxidation state, suggestions about methods to obtain it and definitions of it. To resolve the issue, an IUPAC project (2008-040-1-200) was started in 2008 on the "Comprehensive Definition of Oxidation State", and was concluded by two reports[5][4] and by the revised entries "Oxidation State"[6] and "Oxidation Number"[7] in the IUPAC Gold Book. The outcomes were a single definition of oxidation state and two algorithms to calculate it in molecular and extended-solid compounds, guided by Allen electronegativities that are independent of oxidation state.

See also

References

  1. ^ Wang, G.; Zhou, M.; Goettel, G. T.; Schrobilgen, G. J.; Su, J.; Li, J.; Schlöder, T.; Riedel, S. (2014). "Identification of an iridium-containing compound with a formal oxidation state of IX". Nature. 514: 475–477. doi:10.1038/nature13795.
  2. ^ Yu, H.-S.; Truhlar, D. G. (2016). "Oxidation state 10 exists". Angew. Chem. Int. Ed. 55: 9004–9006. doi:10.1002/anie.201604670.
  3. ^ Lin, J. T.; Hagen, G. P.; Ellis, J. E. (1983). "Highly reduced organometallics. 9. Synthesis and characterization of the tetrasodium tetracarbonylmetalates(4−) of chromium, molybdenum, and tungsten, Na4M(CO)4: their reactions with weak acids to generate H
    2
    M
    2
    (CO)2−
    8
    (M = Cr, Mo, and W)". J. Am. Chem. Soc. 105: 2296–2303. doi:10.1021/ja00346a032.
  4. ^ a b c d Karen, P.; McArdle, P.; Takats, J. (2016). "Comprehensive definition of oxidation state (IUPAC Recommendations 2016)". Pure Appl. Chem. 88: 831–839. doi:10.1515/pac-2015-1204.
  5. ^ a b c d e f g h Karen, P.; McArdle, P.; Takats, J. (2014). "Toward a comprehensive definition of oxidation state (IUPAC Technical Report)". Pure Appl. Chem. 86: 1017–1081. doi:10.1515/pac-2013-0505.
  6. ^ a b c IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "Oxidation state". doi:10.1351/goldbook.O04365
  7. ^ a b IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "Oxidation number". doi:10.1351/goldbook.O04363
  8. ^ Muir, K. W.; Ibers, J. A. (1969). "The structure of chlorocarbonyl(sulfur dioxide)bis(triphenylphosphine)rhodium, RhCl(CO)(SO2)(P(C6H5)3)2". Inorg. Chem. 8: 1921–1928. doi:10.1021/ic50079a024.
  9. ^ Jørgensen, C. K. (1966). "Electric Polarizability, Innocent Ligands and Spectroscopic Oxidation States". Structure and Bonding. 1. Berlin: Springer-Verlag. pp. 234–248.
  10. ^ Karen, P. (2015). "Oxidation state, a long-standing issue!". Angew. Chem. Int. Ed. 54: 4716–4726. doi:10.1002/anie.201407561. PMC 4506524.
  11. ^ Martinie, R. J.; Bultema, J. J.; Wal, M. N. V.; Burkhart, B. J.; Griend, D. A. V.; DeCock, R. L. (2011). "Bond order and chemical properties of BF, CO, and N2". J. Chem. Educ. 88: 1094–1097. doi:10.1021/ed100758t.
  12. ^ a b Petrucci, R. H.; Harwood, W. S.; Herring, F. G. (2002). General Chemistry (8th ed.). Prentice-Hall.
  13. ^ Senn, M. S.; Wright, J. P.; Attfield, J. P. (2012). "Charge order and three-site distortions in the Verwey structure of magnetite". Nature. 481: 173–176. doi:10.1038/nature10704.
  14. ^ Whitten, K. W.; Galley, K. D.; Davis, R. E. (1992). General Chemistry (4th ed.). Saunders. p. 147.
  15. ^ a b c Connelly, N. G.; Damhus, T.; Hartshorn, R. M.; Hutton, A. T. Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005) (PDF). RSC Publishing.
  16. ^ "Antoine Laurent Lavoisier The Chemical Revolution - Landmark - American Chemical Society". American Chemical Society. Retrieved 14 July 2018.
  17. ^ "Lavoisier on Elements". Chem125-oyc.webspace.yale.edu. Retrieved 14 July 2018.
  18. ^ Wöhler, F. (1835). Grundriss der Chemie: Unorganische Chemie [Foundations of Chemistry: Inorganic Chemistry]. Berlin: Duncker und Humblot. p. 4.
  19. ^ Jensen, W. B. (2007). "the origin of the oxidation-state concept". J. Chem. Educ. 84: 1418–1419. doi:10.1021/ed084p1418.
  20. ^ Stock, A. (1919). "Einige Nomenklaturfragen der anorganischen Chemie" [Some nomenclature issues of inorganic chemistry]. Angew. Chem. 32: 373–374. doi:10.1002/ange.19190329802.
  21. ^ a b Jorissen, W. P.; Bassett, H.; Damiens, A.; Fichter, F.; Rémy, H. (1941). "Rules for naming inorganic compounds". J. Am. Chem. Soc. 63: 889–897. doi:10.1021/ja01849a001.
  22. ^ Latimer, W. M. (1938). The Oxidation States of the Elements and their Potentials in Aqueous Solutions (1st ed.). Prentice-Hall.
  23. ^ Bray, W. C.; Branch, G. E. K. (1913). "Valence and tautomerism". J. Am. Chem. Soc. 35: 1440–1447. doi:10.1021/ja02199a003.
  24. ^ Noyes, A. A.; Pitzer, K. S.; Dunn, C. L. (1935). "Argentic salts in acid solution, I. The oxidation and reduction reactions". J. Am. Chem. Soc. 57: 1221–1229. doi:10.1021/ja01310a018.
  25. ^ Noyes, A. A.; Pitzer, K. S.; Dunn, C. L. (1935). "Argentic salts in acid solution, II. The oxidation state of argentic salts". J. Am. Chem. Soc. 57: 1229–1237. doi:10.1021/ja01310a019.
  26. ^ Fernelius, W. C. (1948). "Some problems of inorganic nomenclature". Chem. Eng. News. 26: 161–163. doi:10.1021/cen-v026n003.p161.
  27. ^ Fernelius, W. C.; Larsen, E. M.; Marchi, L. E.; Rollinson, C. L. (1948). "Nomenclature of coördination compounds". Chem. Eng. News. 26: 520–523. doi:10.1021/cen-v026n008.p520.
  28. ^ Pauling, L. (1948). "The modern theory of valency". J. Chem. Soc. 1948: 1461–1467. doi:10.1039/JR9480001461.
  29. ^ Calvert, J. G. (1990). "IUPAC Recommendation 1990". Pure Appl. Chem. 62: 2204. doi:10.1351/pac199062112167.
Acid strength

Acid strength refers to the tendency of an acid, symbolised by the chemical formula HA, to dissociate into a proton, H+, and an anion, A−. The dissociation of a strong acid in solution is effectively complete, except in its most concentrated solutions.

HA → H+ + A−Examples of strong acids are hydrochloric acid (HCl), perchloric acid (HClO4), nitric acid (HNO3) and sulfuric acid (H2SO4).

A weak acid is only partially dissociated, with both the undissociated acid and its dissociation products being present, in solution, in equilibrium with each other.

HA ⇌ H+ + A−.Acetic acid (CH3COOH) is an example of a weak acid. The strength of a weak acid is quantified by its acid dissociation constant, pKa value.

The strength of a weak organic acid may depend on substituent effects. The strength of an inorganic acid is dependent on the oxidation state for the atom to which the proton may be attached. Acid strength is solvent-dependent. For example, hydrogen chloride is a strong acid in aqueous solution, but is a weak acid when dissolved in glacial acetic acid.

Chlorous acid

Chlorous acid is an inorganic compound with the formula HClO2. It is a weak acid. Chlorine has oxidation state +3 in this acid. The pure substance is unstable, disproportionating to hypochlorous acid (Cl oxidation state +1) and chloric acid (Cl oxidation state +5):

2 HClO2 → HClO + HClO3Although the acid is difficult to obtain in pure substance, the conjugate base, chlorite, derived from this acid is stable. One example of a salt of this anion is the well-known sodium chlorite. This and related salts are sometimes used in the production of chlorine dioxide.

Cytochrome

Cytochromes are proteins containing heme as a cofactor. They are classified according to the type of heme and its mode of binding. Four varieties are recognized by the IUBMB, cytochromes a, cytochromes b, cytochromes c and cytochrome d. Cytochrome function is linked to the reversible redox change from ferrous (Fe(II)) to the ferric (Fe(III)) oxidation state of the iron found in the heme core. In addition to the classification by the IUBMB into four cytochrome classes, several additional classifications such as cytochrome o and cytochrome P450 can be found in biochemical literature.

Disproportionation

Disproportionation, sometimes called dismutation, is a redox reaction in which a compound of intermediate oxidation state converts to two different compounds, one of higher and one of lower oxidation states. Although not widely accepted, disproportionation is sometimes used to describe any desymmetrizing reaction of the following type: 2 A → A' + A", regardless of any redox process.

Europium

Europium is a chemical element with symbol Eu and atomic number 63. It was isolated in 1901 and is named after the continent of Europe. It is a moderately hard, silvery metal which readily oxidizes in air and water. Being a typical member of the lanthanide series, europium usually assumes the oxidation state +3, but the oxidation state +2 is also common. All europium compounds with oxidation state +2 are slightly reducing. Europium has no significant biological role and is relatively non-toxic compared to other heavy metals. Most applications of europium exploit the phosphorescence of europium compounds. Europium is one of the least abundant elements in the universe; only about 5×10−8% of all matter in the universe is europium.

Ferrous

In chemistry, ferrous (Fe2+), indicates a divalent iron compound (+2 oxidation state), as opposed to ferric, which indicates a trivalent iron compound (+3 oxidation state). This usage has decreased, with current IUPAC nomenclature having names containing the oxidation state in bracketed Roman numerals instead, such as iron(II) oxide for ferrous oxide (FeO), and iron(III) oxide for ferric oxide (Fe2O3).

Outside chemistry, ferrous indicates the presence of iron. The word is derived from the Latin word ferrum ("iron"). Ferrous metals include steel and pig iron (with a carbon content of a few percent) and alloys of iron with other metals (such as stainless steel). Manipulation of atom-to-atom relationships between iron, carbon, and various alloying elements establishes the specific properties of ferrous metals. Non-ferrous metals and alloys do not contain an appreciable amount of iron.Ferric refers to iron-containing materials or compounds. In chemistry the term is reserved for iron with an oxidation number of +3, also denoted iron(III) or Fe3+. On the other hand, ferrous refers to iron with oxidation number of +2, denoted iron(II) or Fe2+.

Iridium tetroxide

Iridium tetroxide (IrO4, Iridium(VIII) oxide) is a binary compound of oxygen and iridium in oxidation state +VIII. This compound was formed by photochemical rearrangement of [(η1-O2)IrO2] in solid argon at a temperature of 6 K (−267 °C, −449 °F). At higher temperatures the oxide is unstable. The detection of the iridium tetroxide cation by infrared photodissociation spectroscopy with formal oxidation state IX has been reported.

Main-group element

In chemistry and atomic physics, the main group is the group of elements whose lightest members are represented by helium, lithium, beryllium, boron, carbon, nitrogen, oxygen, and fluorine as arranged in the periodic table of the elements. The main group includes the elements (except hydrogen, which is sometimes not included) in groups 1 and 2 (s-block), and groups 13 to 18 (p-block). The s-block elements are primarily characterised by one main oxidation state, and the p-block elements, when they have multiple oxidation states, often have common oxidation states separated by two units.

Main-group elements (with some of the lighter transition metals) are the most abundant elements on earth, in the solar system, and in the universe. They are sometimes also called the representative elements.

Group 12 elements are often considered to be transition metals; however, zinc (Zn), cadmium (Cd), and mercury (Hg) share some properties of both groups, and many scientists believe they should be included in the main group. Occasionally, even the group 3 elements as well as the lanthanides and actinides have been included, because especially the group 3 elements and lanthanides are electropositive elements with only one main oxidation state like the group 1 and 2 elements. The position of the actinides is more questionable, but the most common and stable of them, thorium (Th) and uranium (U), are similar to main-group elements as thorium is an electropositive element with only one main oxidation state (+4), and uranium has two main ones separated by two oxidation units (+4 and +6).In older nomenclature the main-group elements are groups IA and IIA, and groups IIIB to 0 (CAS groups IIIA to VIIIA). Group 12 is labelled as group IIB in both systems. Group 3 is labelled as group IIIA in the older nomenclature (CAS group IIIB).

Oxidative addition

Oxidative addition and reductive elimination are two important and related classes of reactions in organometallic chemistry. Oxidative addition is a process that increases both the oxidation state and coordination number of a metal centre. Oxidative addition is often a step in catalytic cycles, in conjunction with its reverse reaction, reductive elimination.

Permanganate

A permanganate is the general name for a chemical compound containing the manganate(VII) ion, (MnO−4). Because manganese is in the +7 oxidation state, the permanganate(VII) ion is a strong oxidizing agent. The ion has tetrahedral geometry. Permanganate solutions are purple in color and are stable in neutral or slightly alkaline media. The exact chemical reaction is dependent upon the organic contaminants present and the oxidant utilized. For example, trichloroethene (C2H3Cl3) is oxidized by sodium permanganate to form carbon dioxide (CO2), manganese dioxide (MnO2), sodium ions (Na+), hydronium ions (H+), and chloride ions (Cl−).In an acidic solution, permanganate(VII) is reduced to the pale pink +2 oxidation state of the manganese(II) (Mn2+) ion.

8 H+ + MnO−4 + 5 e− → Mn2+ + 4 H2OIn a strongly basic solution, permanganate(VII) is reduced to the green +6 oxidation state of the manganate ion, MnO2−4.

MnO−4 + e− → MnO2−4In a neutral medium, however, it gets reduced to the brown +4 oxidation state of manganese dioxide MnO2.

2 H2O + MnO−4 + 3 e− → MnO2 + 4 OH−

Peroxide

Peroxides are a group of compounds with the structure R−O−O−R. The O−O group in a peroxide is called the peroxide group or peroxo group. In contrast to oxide ions, the oxygen atoms in the peroxide ion have an oxidation state of −1.

The most common peroxide is hydrogen peroxide (H2O2), colloquially known simply as "peroxide". It is marketed as a solution in water at various concentrations. Since hydrogen peroxide is colorless, so are these solutions. It is mainly used as an oxidant and bleaching agent. However, hydrogen peroxide is also biochemically produced in the human body, largely as a result of a range of oxidase enzymes. Concentrated solutions are potentially dangerous when in contact with organic compounds.

Aside from hydrogen peroxide, some other major classes of peroxides are these:

Peroxy acids, the peroxy derivatives of many familiar acids, examples being peroxymonosulfuric acid and peracetic acid.

Metal peroxides, examples being barium peroxide (BaO2) and sodium peroxide (Na2O2).

Organic peroxides, compounds with the linkage C−O−O−C or C−O−O−H. One example is tert-butylhydroperoxide

Main group peroxides, compounds with the linkage E−O−O−E (E = main group element), one example is potassium peroxydisulfate.

Potassium ferrate

Potassium ferrate is the chemical compound with the formula K2FeO4. This purple salt is paramagnetic, and is a rare example of an iron(VI) compound. In most of its compounds, iron has the oxidation state +2 or +3 (Fe2+ or Fe3+). Reflecting its high oxidation state, FeO42− is a powerful oxidizing agent.

Praseodymium

Praseodymium is a chemical element with symbol Pr and atomic number 59. It is the third member of the lanthanide series and is traditionally considered to be one of the rare-earth metals. Praseodymium is a soft, silvery, malleable and ductile metal, valued for its magnetic, electrical, chemical, and optical properties. It is too reactive to be found in native form, and pure praseodymium metal slowly develops a green oxide coating when exposed to air.

Praseodymium always occurs naturally together with the other rare-earth metals. It is the fourth most common rare-earth element, making up 9.1 parts per million of the Earth's crust, an abundance similar to that of boron. In 1841, Swedish chemist Carl Gustav Mosander extracted a rare-earth oxide residue he called didymium from a residue he called "lanthana", in turn separated from cerium salts. In 1885, the Austrian chemist Baron Carl Auer von Welsbach separated didymium into two elements that gave salts of different colours, which he named praseodymium and neodymium. The name praseodymium comes from the Greek prasinos (πράσινος), meaning "green", and didymos (δίδυμος), "twin".

Like most rare-earth elements, praseodymium most readily forms the +3 oxidation state, which is the only stable state in aqueous solution, although the +4 oxidation state is known in some solid compounds and, uniquely among the lanthanides, the +5 oxidation state is attainable in matrix-isolation conditions. Aqueous praseodymium ions are yellowish-green, and similarly praseodymium results in various shades of yellow-green when incorporated into glasses. Many of praseodymium's industrial uses involve its ability to filter yellow light from light sources.

Redox

Redox (short for reduction–oxidation reaction) (pronunciation: redoks or reedoks) is a chemical reaction in which the oxidation states of atoms are changed. Any such reaction involves both a reduction process and a complementary oxidation process, two key concepts involved with electron transfer processes. Redox reactions include all chemical reactions in which atoms have their oxidation state changed; in general, redox reactions involve the transfer of electrons between chemical species. The chemical species from which the electron is stripped is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. It can be explained in simple terms:

Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion.

Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.As an example, during the combustion of wood, oxygen from the air is reduced, gaining electrons from carbon which is oxidized. Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, oxygen is not necessarily included in such reactions, as other chemical species can serve the same function.The reaction can occur relatively slowly, as with the formation of rust, or more quickly, in the case of fire. There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide (CO2) or the reduction of carbon by hydrogen to yield methane (CH4), and more complex processes such as the oxidation of glucose (C6H12O6) in the human body.

Thallium halides

The thallium halides include monohalides, where thallium has oxidation state +1, trihalides where thallium generally has oxidation state +3 and some intermediate halides with mixed +1 and +3 oxidation states. These materials find use in specialized optical settings, such as focusing elements in research spectrophotometers. Compared to the more common zinc selenide-based optics, materials such as thallium bromoiodide enable transmission at longer wavelengths. In the infrared, this allows for measurements as low as 350 cm−1 (28 µm), whereas zinc selenide is opaque by 21.5 µm and ZnSe optics are generally only usable to 650 cm−1 (15 µm).

Transition metal

In chemistry, the term transition metal (or transition element) has three possible meanings:

The IUPAC definition defines a transition metal as "an element whose atom has a partially filled d sub-shell, or which can give rise to cations with an incomplete d sub-shell".

Many scientists describe a "transition metal" as any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table. In actual practice, the f-block lanthanide and actinide series are also considered transition metals and are called "inner transition metals".

Cotton and Wilkinson expand the brief IUPAC definition (see above) by specifying which elements are included. As well as the elements of groups 4 to 11, they add scandium and yttrium in group 3, which have a partially filled d subshell in the metallic state. Lanthanum and actinium in group 3 are, however, classified as lanthanides and actinides respectively.English chemist Charles Bury (1890–1968) first used the word transition in this context in 1921, when he referred to a transition series of elements during the change of an inner layer of electrons (for example n = 3 in the 4th row of the periodic table) from a stable group of 8 to one of 18, or from 18 to 32. These elements are now known as the d-block.

Valence (chemistry)

In chemistry, the valence or valency of an element is a measure of its combining power with other atoms when it forms chemical compounds or molecules. The concept of valence developed in the second half of the 19th century and helped successfully explain the molecular structure of inorganic and organic compounds.

The quest for the underlying causes of valence led to the modern theories of chemical bonding, including the cubical atom (1902), Lewis structures (1916), valence bond theory (1927), molecular orbitals (1928), valence shell electron pair repulsion theory (1958), and all of the advanced methods of quantum chemistry.

Xenon tetroxide

Xenon tetroxide is a chemical compound of xenon and oxygen with molecular formula XeO4, remarkable for being a relatively stable compound of a noble gas. It is a yellow crystalline solid that is stable below −35.9 °C; above that temperature it is very prone to exploding and decomposing into elemental xenon and oxygen (O2).All eight valence electrons of xenon are involved in the bonds with the oxygen, and the oxidation state of the xenon atom is +8. Oxygen is the only element that can bring xenon up to its highest oxidation state; even fluorine can only give XeF6 (+6).

Two other short-lived xenon compounds with an oxidation state of +8, XeO3F2 and XeO2F4, are accessible by the reaction of xenon tetroxide with xenon hexafluoride. XeO3F2 and XeO2F4 can be detected with mass spectrometry. The perxenates are also compounds where xenon has the +8 oxidation state.

Ytterbium

Ytterbium is a chemical element with symbol Yb and atomic number 70. It is the fourteenth and penultimate element in the lanthanide series, which is the basis of the relative stability of its +2 oxidation state. However, like the other lanthanides, its most common oxidation state is +3, as in its oxide, halides, and other compounds. In aqueous solution, like compounds of other late lanthanides, soluble ytterbium compounds form complexes with nine water molecules. Because of its closed-shell electron configuration, its density and melting and boiling points differ significantly from those of most other lanthanides.

In 1878, the Swiss chemist Jean Charles Galissard de Marignac separated from the rare earth "erbia" another independent component, which he called "ytterbia", for Ytterby, the village in Sweden near where he found the new component of erbium. He suspected that ytterbia was a compound of a new element that he called "ytterbium" (in total, four elements were named after the village, the others being yttrium, terbium and erbium). In 1907, the new earth "lutecia" was separated from ytterbia, from which the element "lutecium" (now lutetium) was extracted by Georges Urbain, Carl Auer von Welsbach, and Charles James. After some discussion, Marignac's name "ytterbium" was retained. A relatively pure sample of the metal was not obtained until 1953. At present, ytterbium is mainly used as a dopant of stainless steel or active laser media, and less often as a gamma ray source.

Natural ytterbium is a mixture of seven stable isotopes, which altogether are present at concentrations of 3 parts per million. This element is mined in China, the United States, Brazil, and India in form of the minerals monazite, euxenite, and xenotime. The ytterbium concentration is low because it is found only among many other rare earth elements; moreover, it is among the least abundant. Once extracted and prepared, ytterbium is somewhat hazardous as an eye and skin irritant. The metal is a fire and explosion hazard.

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