Oxidation state
The oxidation state is an indicator of the degree of oxidation of an atom in a chemical compound. The formal oxidation state is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic. Oxidation states are typically represented by integers, which can be positive, negative, or zero. In some cases, the average oxidation state of an element is a fraction, such as 8/3 for iron in magnetite (Fe
3O
4). The highest known oxidation state is +8 in the tetroxides (O4) of ruthenium, xenon, osmium, iridium, and hassium, and some complexes involving plutonium, while the lowest known oxidation state is −4 for some elements in the carbon group.
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 definition of the oxidation state listed by IUPAC is as follows:[1]
[Oxidation state] is defined as the charge an atom might be imagined to have when electrons are counted according to an agreed-upon set of rules:
- the oxidation state of a free element (uncombined element) is zero
- for a simple (monoatomic) ion, the oxidation state is equal to the net charge on the ion
- hydrogen has an oxidation state of 1 and oxygen has an oxidation state of −2 when they are present in most compounds. (Exceptions to this are that hydrogen has an oxidation state of −1 in hydrides of active metals, e.g. LiH, and oxygen has an oxidation state of −1 in peroxides, e.g. H2O2
- the algebraic sum of oxidation states of all atoms in a neutral molecule must be zero, while in ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion.
In the nomenclature of inorganic compounds, the oxidation state (or number) is represented by a Roman numeral. In older literature the term is referred to as Stock number, however the use of this term is no longer recommended by IUPAC.[2] The oxidation state in compound naming is placed either as a right superscript to the element symbol in chemical formula, for example FeIII, or in parentheses after the name of the element, iron(III) in chemical names. For example Fe2(SO4)3 is named iron(III) sulfate and its formula can be shown as FeIII2(SO4)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.[3]
General rules for determining oxidation states without use of Lewis structures
There are general rules (adopted by IUPAC in 1990) for determining the oxidation states of atoms in simple chemical compounds without the use of structural formulae:[1]
- Any pure element—even if it forms diatomic molecules like chlorine (Cl2)—has an oxidation state (OS) of zero. Examples of this are Cu or O2.
- For monatomic ions, the OS is the same as the charge of the ion. For example, the sulfide anion (S2−) has an OS of −2, whereas the lithium cation (Li+) has an OS of +1.
- The sum of oxidation states for all atoms in a molecule or polyatomic ion is equal to the charge of the molecule or ion. Thus, the OS of one element can be calculated from the oxidation states of the other elements.
- An application of this rule is that the sum of the oxidation states of all atoms in a neutral molecule must be zero. Consider a neutral molecule of carbon dioxide, CO2. Oxygen is assumed to have its usual oxidation state of −2, and so the sum of the oxidation states of all the atoms can be expressed as OS(C) + 2(−2) = 0, or OS(C) − 4 = 0, where OS(C) is the unknown oxidation state of carbon. Thus, it can be seen that the oxidation state of carbon in the molecule is +4.
- Likewise, in polyatomic ions, the sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. As an example, consider the sulfite anion, which has the formula SO32-. As indicated by the formula, the total charge of this ion is −2. Because all three oxygen atoms are assumed to have their usual oxidation state of −2, and the sum of the oxidation states of the all the atoms is equal to the charge of the ion, the sum of the oxidation states can be represented as OS(S) + 3(−2) = −2, or OS(S) − 6 = −2, where OS(S) is the unknown oxidation state of sulfur. Thus, it can be computed that OS(S) = +4.
These facts, combined with the fact that some elements almost always have certain oxidation states (due to their very high electropositivity or electronegativity), allows one to compute the oxidation states for the remaining atoms (such as transition metals) in simple compounds.
The following rules can be used for initially assigning oxidation states for certain elements, in simple compounds:
- Fluorine has an oxidation state of −1 when bonded to any other element, since it has the highest electronegativity of all reactive elements.
- Halogens other than fluorine have an oxidation state of −1 except when they are bonded to oxygen, to nitrogen, or to another halogen which is more electronegative. For example, the oxidation state of chlorine in chlorine monofluoride (ClF) is +1. However, in bromine monochloride (BrCl), the oxidation state of Cl is −1.
- Hydrogen has an oxidation state of +1 except when bonded to more electropositive elements such as sodium, aluminium, and boron, as in NaH, NaBH
4, LiAlH
4, where each H has an oxidation state of −1. - In compounds, oxygen typically has an oxidation state of −2, though there are exceptions that are listed below, such as peroxides (e.g. hydrogen peroxide H2O2), where oxygen has an OS of −1.
- Alkali metals have an oxidation state of +1 in virtually all of their compounds (exception, see alkalide).
- Alkaline earth metals have an oxidation state of +2 in virtually all of their compounds.
Example for a complex salt: In Cr(OH)
3, oxygen has an oxidation state of −2 (no fluorine or O–O bonds present), and hydrogen has a state of +1 (bonded to oxygen). So, each of the three hydroxide groups has an oxidation state of −2 + 1 = −1. As the compound is neutral, chromium has an oxidation state of +3.
Using electronegativity
In the 1970[4] rules IUPAC recommended that oxidation state was used in nomenclature and elsewhere in inorganic chemistry as the "charge that would be present on an atom if the electrons were assigned to the more electronegative atom", but with a convention that hydrogen is considered to be positive in combination with non-metals and a bond between like atoms makes no contribution to the oxidation number.
The use of electronegativity in this way was introduced by Pauling in 1947.[5] This method of determining oxidation state is found in some recent text books (example see reference [6]). This method allows the oxidation state of all atoms in a molecule to be determined whereas the IUPAC 1990/2005 definition does not.[7]
Oxidation state and formal charge
The oxidation state of an atom is often different from the formal charge often included in Lewis structures (when it is non-zero). The oxidation state is calculated by assuming that each chemical bond (except between identical atoms) is ionic so that both electrons are assigned to the more electronegative bonded atom. In contrast, the formal charge is calculated by assuming that each bonds is covalent so that one electron is assigned to each bonded atom. For example, in ammonium ion (NH4+) the oxidation state of nitrogen is -3, as all eight valence electrons are assigned to the nitrogen atom which is more electronegative than hydrogen. However the formal charge is +1, calculated by assigning only four valence electrons (one per bond) to nitrogen. For comparison, the nitrogen in ammonia (NH3) has oxidation state -3 also but a formal charge of zero. On protonation of ammonia the formal charge on nitrogen changes, but its oxidation state does not for molecules which contain nonequivalent atoms of the same element.
Calculation of formal oxidation states with a Lewis structure
There are two common ways of computing the oxidation state of an atom in a compound. The first is the simple algebraic sum technique above, used in compounds that do not require a Lewis structure. The second is used for molecules when one has a Lewis structure.
It should be remembered that the oxidation state of an atom does not represent the "real" charge on that atom: This is particularly true of high oxidation states, where the ionization energy required to produce a multiply positive ion are far greater than the energies available in chemical reactions. The assignment of electrons between atoms in calculating an oxidation state is purely a formalism, but is a useful one for the understanding of many chemical reactions.
For more about issues with calculating atomic charges, see partial charge.
The Lewis structure
When a Lewis structure of a molecule is available, the oxidation states may be assigned by computing the difference between the number of valence electrons that a neutral atom of that element would have and the number of electrons that "belong" to it in the Lewis structure. For purposes of computing oxidation states, electrons in a bond between atoms of different elements belong to the more electronegative atom; electrons in a bond between atoms of the same element are split equally, and electrons in a lone pair belong only to the atom with the lone pair.
For example, consider acetic acid:
The methyl group carbon atom has 6 valence electrons from its bonds to the hydrogen atoms because carbon is more electronegative than hydrogen. Also, 1 electron is gained from its bond with the other carbon atom because the electron pair in the C–C bond is split equally, giving a total of 7 electrons. A neutral carbon atom would have 4 valence electrons, because carbon is in group 14 of the periodic table. The difference, 4 – 7 = –3, is the oxidation state of that carbon atom. That is, if it is assumed that all the bonds were 100% ionic (which in fact they are not), the carbon would be described as C3-.
Following the same rules, the carboxylic acid carbon atom has an oxidation state of +3 (it only gets one valence electron from the C–C bond; the oxygen atoms get all the other electrons because oxygen is more electronegative than carbon). The oxygen atoms both have an oxidation state of –2; they get 8 electrons each (4 from the lone pairs and 4 from the bonds), while a neutral oxygen atom would have 6. The hydrogen atoms all have oxidation state +1, because they surrender their electron to the more electronegative atoms to which they are bonded.
Inequivalent atoms of an element
An example of a molecule with inequivalent atoms of the same element is the thiosulfate ion (S2O32−), for which the algebraic sum rule yields the average value +2 for sulfur, where the two ionizing electrons are assigned to the terminal sulfur atom. However, the use of a Lewis structure and electron counting shows that the two sulfur atoms are different. The central sulfur is assigned only one valence electron from the S-S bond and no valence electrons from the S-O bonds, compared to six valence electrons for a free sulfur atom, so the oxidation state of the central sulfur is +5. The terminal sulfur atom is assigned the other electron from the S-S bond plus three lone pairs for a total of seven valence electrons, so its oxidation state is −1.
Redox reactions
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 the Tollens' reagent to acetic acid (shown below), the carbonyl carbon atom changes its oxidation state from +1 to +3 (oxidation). This oxidation is balanced by reducing two equivalents of silver from Ag+ to Ag0.
In such structural diagrams for organic chemistry, oxidation states are represented by Roman numerals to distinguish them from formal charges (calculated with all bonds covalent).
Elements with multiple oxidation states
Most elements have more than one possible oxidation state. For example, carbon has nine integer oxidation states:
| Oxidation state | Example compound |
|---|---|
| –4 | CH 4 |
| –3 | C 2H 6 |
| –2 | CH 3Cl |
| –1 | C 2H 2 |
| 0 | CH 2Cl 2 |
| +1 | CHCl 2—CHCl 2 |
| +2 | CHCl 3 |
| +3 | C 2Cl 6 |
| +4 | CCl 4 |
Fractional oxidation states
Fractional oxidation states are often used to represent the average oxidation states of several atoms in a structure. For example, in KO
2, the diatomic superoxide ion has an overall charge of −1, so each of its two oxygen atoms is assigned an oxidation state of −½, This ion is 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 ion C
5H
5−, the oxidation state of C is (−1) + (−1⁄5) = −6⁄5. The −1 occurs because each C is bonded to one hydrogen (a less electronegative element), and the −1⁄5 because the total ionic charge is divided among five equivalent C.
If the average refers to atoms that are not equivalent, the average oxidation state may not be representative of each of the atoms. This is true in magnetite Fe
3O
4, whose formula leads to an average oxidation state of +8⁄3. In fact, two-thirds of the iron ions are Fe3+, and one-third Fe2+.
Likewise, the ozonide ion O3− has an average oxidation state of −1⁄3. However, this ion is V-shaped, meaning that the central oxygen is not equivalent to the two others and cannot be assumed to have the same oxidation state.
As an example, some species contain carbon in more than one oxidation state, giving a fractional oxidation state overall:
| Oxidation state | Example species |
|---|---|
| –6⁄5 | C 5H 5− |
| –6⁄7 | C 7H 7+ |
| –5⁄4 | C 8H 82− |
Oxidation number in coordination compounds
The terms oxidation state and oxidation number are often used interchangeably. However, oxidation number is used in coordination chemistry with a slightly different meaning. In coordination chemistry, the rules used for counting electrons are different: Every electron in a metal-ligand bond belongs to the ligand, regardless of electronegativity, so that the oxidation number is the charge that would remain if all ligands were removed together with the electron pairs shared with the central atom.[2]
For most coordination complexes, the metal atom is the less electronegative end of each metal-ligand bond, so that this rule gives the same result as the electronegativity-based rule[7] There are exceptions, however, such as Wilkinson's catalyst RhCl(PPh3)3 (Ph = phenyl), in which the rhodium atom is more electronegative than phosphorus. Nevertheless the oxidation number of rhodium in this molecule is considered to be +1 and the molecule’s systematic name is chlorotris(triphenylphosphine)rhodium(I), as the Rh-P electrons are assigned to the P atom of the ligand. The electronegativity rule would assign them instead to the Rh with an oxidation state of -5.
Spectroscopic oxidation states vs. formal oxidation numbers
Although formal oxidation numbers can be helpful for classifying compounds, they are unmeasurable and their physical meaning can be ambiguous. Formal oxidation numbers require particular caution for molecules where the bonding is covalent, since the formal oxidation numbers require the heterolytic removal of ligands, which essentially denies covalency. Spectroscopic oxidation states, as defined by Jorgenson and reiterated by Wieghardt, are measurables that are bench-marked using spectroscopic and crystallographic data.[8]
Oxidation state can also have effect on spectroscopic studies of compounds. In infrared spectroscopy of metal carbonyls this effect is illustrated by using spectroscopic studies on metals from oxidation states of –2 to +2.
Unusual formal oxidation states
Unusual formal oxidation states of metals are important in biochemical processes, the notable ones being Fe(IV) and Fe(V) in Cytochrome P450-containing systems.
History of the oxidation number concept
Oxidation itself was first studied by Antoine Lavoisier, who believed that oxidation was always the result of reactions with oxygen,[9] thus the name. Although Lavoisier's idea has been shown to be incorrect, the name he proposed is still used, albeit more generally.
Oxidation states were one of the intellectual "stepping stones" that Mendeleev used to derive the periodic table.
The Stock nomenclature (named for Alfred Stock who suggested it in 1919) was intended to replace the naming that was prevalent at the time. Under the Stock system FeCl2 was called iron(II) chloride rather than ferrous chloride.
The current concept of "oxidation state" was introduced by W. M. Latimer in 1938.[10] In 1940 IUPAC recommended that the term Stock number should be replaced by the term oxidation number. In 1947 Pauling proposed that the oxidation number could be determined using the electronegativity of the atoms to determine the "ions" in the formal determination of oxidation number.[5] In 1970[4] IUPAC defined oxidation number in terms of electronegativity. In 1990 IUPAC changed course and adopted a rule based determination for the "central atom" rather than using electronegativity. This is the definition in the current gold book for "oxidation state". They also introduced the definition of oxidation number, shown in the current gold book, that appears to make oxidation number specific to coordination chemistry. This may not have been their intention as in 2005 they issued new recommendations for inorganic nomenclature that define oxidation number in the same terms as the 1990 definition of oxidation state, and that, oxidation number is, as in the earlier recommendations, used in the naming of inorganic compounds.
Oxidation number versus oxidation state
In the wider field of chemistry these definitions have not generally been adhered to and both terms are used interchangeably, as they were when Latimer introduced the concept in 1938.[10] For example two well known text books[11][12] use the term oxidation state and represent it in Roman numerals in chemical formulae. The point has been made that if there is any semantic difference between the terms, then oxidation number refers to the specific numerical value assigned to the entity known as oxidation state, much as IUPAC use the term charge number to refer to the numerical value assigned to the entity know as ionic charge.[13] The IUPAC Gold Book takes the definitions from 1990 IUPAC papers[14][15] rather than the more recent current IUPAC 2005 recommendations. There is a current IUPAC project, "Towards a comprehensive definition of oxidation state", (project 2008-040-1-200) started in 2009 which has yet to report (March 2013). The project was undertaken because the current definition in the IUPAC Gold Book was seen to be "narrow and circular", and "inapplicable to clusters, Zintl phases and some organometallic complexes".
See also
References
- ^ a b IUPAC Gold Book definition: oxidation state PDF
- ^ a b IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "oxidation number". doi:10.1351/goldbook.O04363
- ^ Nomenclature of Inorganic Chemistry IUPAC Recommendations 2005 ed. N. G. Connelly et al. RSC Publishing http://www.chem.qmul.ac.uk/iupac/bioinorg/
- ^ a b Nomenclature of Inorganic chemistry, 2d Edition, Definitive rules 1970, Butterworths
- ^ a b General Chemistry: An Introduction to Descriptive Chemistry and Modern Chemical Theory, Linus Pauling, W.H Freeman, 1947
- ^ Basic Concepts of Chemistry, 8th Edition, Leo J. Malone, Theodore Dolter, John Wiley & Sons, 2008, ISBN 047174154X , ISBN 978-0471741541
- ^ a b Loock, Hans-Peter (2011). "Expanded Definition of the Oxidation State". Journal of Chemical Education. 88 (3): 282–283. doi:10.1021/ed1005213. ISSN 0021-9584.
- ^ Bill, E. (2005). "Molecular and electronic structure of four- and five-coordinate cobalt complexes containing two o-phenylenediamine- or two o-aminophenol-type ligands at various oxidation levels functional, and correlated ab initio study". Chemistry - A European Journal. 11 (1): 204–224. doi:10.1002/chem.200400850. PMID 15549762.
{{cite journal}}: Unknown parameter|coauthors=ignored (|author=suggested) (help) - ^ The Origin of the Oxidation-State Concept William B. Jensen J. Chem. Educ. 2007, 84, 1418
- ^ a b Jensen, William B. (2007). "The Origin of the Oxidation-State Concept". Journal of Chemical Education. 84 (9): 1418. doi:10.1021/ed084p1418. ISSN 0021-9584.
- ^ Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999), Advanced Inorganic Chemistry (6th ed.), New York: Wiley-Interscience, ISBN 0-471-19957-5
- ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
- ^ Jensen, William B. (2011). "Oxidation States versus Oxidation Numbers". Journal of Chemical Education. 88 (12): 1599–1600. doi:10.1021/ed2001347. ISSN 0021-9584.
- ^ Red Book: IUPAC Nomenclature of Inorganic Chemistry. Third Edition, Blackwell Scientific Publications, Oxford, 1990.
- ^ Calvert, J. G. (1990). "Glossary of atmospheric chemistry terms (Recommendations 1990)". Pure and Applied Chemistry. 62 (11): 2167–2219. doi:10.1351/pac199062112167. ISSN 0033-4545.