What is the oxidation state for iron

Oxidation number

The Oxidation number (also Oxidation state, Oxidation value, electrochemical valence) indicates the ionic charge of an atom within a chemical compound or polyatomic ion that would be present if the compound or polyatomic ion were made up of monatomic ions.[1][2] For this purpose, binding electron pairs are mentally assigned to the more electronegative binding partner. Bonding electron pairs between the same atoms are shared. The atoms in modifications of the elements have an oxidation number of zero. In the case of monatomic ions, the oxidation number is equal to the actual charge. An oxidation number is a numerical value with a sign (+/-).

The oxidation number is a formal parameter and often has little to do with a real charge. However, it is an important formalism for the stoichiometry of redox reactions.

According to IUPAC, the designations oxidation state[1] and oxidation number[3] used. The designation oxidation state corresponds to the oxidation number described here. The oxidation number is an element of the nomenclature of inorganic salts (e.g. for iron (III) chloride), complex compounds (e.g. for potassium hexacyanidoferrate (II)) and complex ions. In the nomenclature of compounds (oxidation number) only whole-number oxidation numbers are used and, in addition to the number 0, only given in Roman numerals. In the case of complex compounds, the value indicates the oxidation number of the central atom.

Another definition reads: The oxidation number of an atom in a chemical compound is formally a measure for specifying the proportions of the electron density around this atom. A positive oxidation number indicates that the electron density is reduced compared to its normal state, a negative one indicates that the electron density around the atom has increased.

Use

The oxidation numbers are used in redox reactions to better recognize the processes. A reduction in the oxidation number of an element through a redox reaction means that this element has been reduced; analogously, an increase in the oxidation number of an element means that it has been oxidized.

Specification of the oxidation number

The oxidation numbers are usually given as Arabic numerals to represent formulas[4] and are usually whole numbers. In contrast to ion charges, the signs are plus and minus in front the numbers indicated. To represent redox reactions, the oxidation numbers are above an element symbol. Occasionally, however, Roman numerals are also used. The possible oxidation numbers of the chemical elements are listed here.

$ \ mathrm {{\ overset {+1} {K}} {\ overset {+7} {Mn}} {\ overset {-2} {O_4}}} $$ \ mathrm {{\ overset {+4} {Mn}} {\ overset {-2} {O_2}}} $$ \ mathrm {{\ overset {+6} {S}} {\ overset {-2} {O_4 ^ {2-}}}} $$ \ mathrm {{\ overset {+4} {S}} {\ overset {-2} {O_3 ^ {2-}}}} $$ \ mathrm {{\ overset {-3} {N}} {\ overset {+1} {H_3}}} $$ \ mathrm {{\ overset {-3} {N}} {\ overset {+1} {H_4 ^ +}}} $$ \ mathrm {{\ overset {+1} {H_2}} {\ overset {-2} {S}}} $$ \ mathrm {{\ overset {0} {O}} \ mathord = \ mathord {\ overset {0} {O}}} $$ \ mathrm {{\ overset {+2} {Fe ^ {2+}}}} $
Potassium perm
manganate
Manganese-
dioxide
Sulfate ionSulfite ionammoniaammonium
-Ion
Sulfur-
hydrogen
oxygenIron (II) ion

When formulating redox reactions, often only the oxidation numbers of the elements that are decisive for the reaction are given:

$ \ mathrm {{\ overset {+7} {Mn}} O_4 ^ - \ + \ 8 \ H_3O ^ + + 5 \ e ^ - \ longrightarrow {\ overset {+2} {Mn ^ {2+}}} + 12 \ H_2O} $
Partial reaction of a redox reaction: reduction of the oxidizing agent permanganate.

The oxidation numbers can also assume fractional values. In the case of hyperoxides such as potassium hyperoxide (KO2), the oxygen atoms have the oxidation number −0.5 and differ from the peroxides with the oxidation number −1.

In the Fe3O4 (Iron (II, III) oxide) Iron has an average oxidation number of +8/3. The oxidation states listed in Roman numerals in the name indicate that the iron atoms in this compound have the oxidation states +2 and +3 are present. (It does not have an oxidation state of +2.3.) FeIIFe2IIIO4 has an inverse spinel structure (simplified: FeO · Fe2O3) and formal Fe2+- and Fe3+-Ions can be localized.

The thiosulfate ion (p2O32−, Disulfate (II) ion) consists of two dissimilar sulfur atoms. The mean oxidation state of sulfur is +2. The discrete levels that can be derived from the structure are +5 and −1. The medium and discrete oxidation states are suitable for stoichiometric calculations in inorganic chemistry.

$ \ mathrm {{\ overset {+1} {K_2}} {\ overset {-1} {O_2}}} $$ \ mathrm {{\ overset {+1} {K}} {\ overset {-0.5} {O_2}}} $$ \ mathrm {{\ overset {+2} {Fe}} (II) {\ overset {+3} {Fe_2}} (III) {\ overset {-2} {O_4}}} $$ \ mathrm {{\ overset {+ \ frac {8} {3}} {Fe_3}} {\ overset {-2} {O_4}}} $$ \ mathrm {{\ overset {+2} {S_2}} {\ overset {-2} {O_3 ^ {2-}}}} $$ \ mathrm {({\ overset {-2} {O_3}} {\ overset {+5} {S}} \ mathord- {\ overset {-1} {S}) ^ {2-}}} $
Potassium peroxidePotassium hyperoxideIron (II, III) oxide
with discrete oxidation states
Iron (II, III) oxide
with a medium oxidation state
Thiosulfate ion
Disulfate (II) ion
with a medium oxidation state
Thiosulfate ion
with discrete oxidation states

In the case of organic compounds, the oxidation numbers are determined separately for each carbon atom:

$ \ mathrm {{\ overset {-4} {C}} H_4} $$ \ mathrm {{\ overset {-3} {C}} H_3 \ mathord- {\ overset {-3} {C}} H_3} $$ \ mathrm {{\ overset {-3} {C}} H_3 \ mathord- {\ overset {-2} {C}} H_2 \ mathord- {\ overset {-3} {C}} H_3} $$ \ mathrm {{\ overset {-3} {C}} H_3 \ mathord- {\ overset {-1} {C}} H \ mathord = {\ overset {-2} {C}} H_2} $
methaneEthanepropanePropene

By comparing oxidation numbers, it can be seen, for example, that a conversion of a primary alcohol to an aldehyde or the conversion of an aldehyde to a carboxylic acid are oxidations.

$ \ mathrm {{\ overset {-3} {C}} H_3 \ mathord- {\ overset {-1} {C}} H_2OH} $$ \ mathrm {{\ overset {-3} {C}} H_3 \ mathord- {\ overset {+1} {C}} HO} $$ \ mathrm {{\ overset {-3} {C}} H_3 \ mathord- {\ overset {+3} {C}} OOH} $
Ethanolacetaldehydeacetic acid

In the case of organic reactions, simple stoichiometric considerations are usually insufficient to describe the reaction. Therefore, compared to inorganic chemistry, oxidation states play a subordinate role. A stoichiometric reaction is such. B. the great test:

$ \ mathrm {CH_3 \ mathord- {\ overset {+1} {C}} HO \ + \ 2 \ {\ overset {+1} {Ag ^ +}} + 2 \ OH ^ - \ longrightarrow \ CH_3 \ mathord - {\ overset {+3} {C}} OOH \ + \ 2 \ {\ overset {0} {Ag}} \ + \ 2 \ H_2O} $
Oxidation of acetaldehyde to acetic acid

In principle, the sum of the oxidation numbers of the atoms in a molecular compound is zero. In the case of ions, the sum of the oxidation numbers is equal to the ion charge. In the case of redox reactions that are stoichiometrically correct, the sum of the oxidation numbers of the starting materials is equal to the sum of the oxidation numbers of the products.

Determination of the oxidation number

Main rules

The oxidation number can be derived using the following rules:

  1. Atoms in the elementary state always have the oxidation number 0 (e.g. I.2, C, O2, P4, S.8; 0 is also possible in connection with other elements).
  2. In the case of monatomic ions, the oxidation number corresponds to the ion charge (e.g. Cu2+ has the oxidation number +2, Ag+ has the oxidation number +1).
  3. The sum of the oxidation numbers of all atoms of a polyatomic neutral compound is equal to 0.
  4. The sum of the oxidation numbers of all atoms of a polyatomic ion is equal to the total charge of this ion.
  5. In the case of covalently formulated compounds (so-called valence line formulas, Lewis formulas), the connection is formally divided into ions. It is assumed that the electrons involved in a bond are completely taken over by the more electronegative atom.
  6. Most elements can occur in several oxidation states.

Auxiliary rules

In practice, it has proven helpful to formulate a few rules for determining the oxidation number:

  1. The fluorine atom (F) as the element with the highest electronegativity always has the oxidation number −1 in compounds.
  2. Oxygen atoms get the oxidation number −2. With 3 exceptions: In peroxides (then: −1) and in hyperoxides (then −0.5) and in connection with fluorine (then: +2).
  3. Other halogen atoms (such as chlorine, bromine, iodine) generally have the oxidation number (−1), except in connection with oxygen or a halogen that is higher in the periodic table.
  4. In compounds, metal atoms always have a positive oxidation number as ions.
  5. Alkali metals always have +1 and alkaline earth metals always +2 as the oxidation number.
  6. Hydrogen atoms get the oxidation number +1, except when hydrogen is directly linked to more “electropositive” atoms such as metals (hydrides) or to itself.
  7. In ionic compounds (salts), the sum of the oxidation numbers is identical to the ionic charge.
  8. In covalent bonds (molecules) the binding electrons are assigned to the more electronegative binding partner. Identical binding partners each receive half of the binding electrons. The oxidation number thus corresponds to the assigned binding electrons compared to the number of external electrons normally present.
  9. The highest possible oxidation number of an element corresponds to the number of major or minor groups in the periodic table (PSE)

Determination based on electronegativity

Determination of the oxidation numbers using 5-hydroxycytosine as an example

If the molecule has a Lewis formula, the oxidation numbers can easily be determined from the electronegativity of the respective elements. To do this, you mentally split every bond and calculate which atom would then get the binding electrons. This depends on the electronegativity; the atom with the greater electronegativity receives the binding electrons. This changes the charge of the atom, the charge then corresponds to the oxidation number. Splitting the bonds is just a mind game, the bonds are not actually split.

The graphic on the right shows an example of the procedure for determining the oxidation numbers of the atoms of the 5-hydroxycytosine molecule. As an example, the procedure on the carbon atom with the oxidation number ± 0 will now be explained: This carbon atom forms three bonds to neighboring atoms, to nitrogen, hydrogen and a double bond to another carbon atom. Now the electronegativities of these elements are compared; Carbon has an electronegativity of 2.55.

  • Nitrogen has an electronegativity of 3.04. Since this is greater than that of carbon, the nitrogen would get both binding electrons in the event of an imaginary cleavage of the bond.
  • Hydrogen has an electronegativity of 2.2. Since this is smaller than that of carbon, the carbon would get both binding electrons in the event of an imaginary cleavage of the bond.
  • The upper carbon atom, of course, also has an electronegativity of 2.55. Therefore, in the event of an imaginary cleavage of the bond, the two carbon atoms share the bond electrons. Since it is a double bond, both would get two.

So when added, the carbon atom has four binding electrons. Elemental carbon also has four binding electrons, so its charge has not changed due to the imaginary cleavage. Its oxidation number is 0.

In comparison, the lowest nitrogen atom gets six bonding electrons in the case of an imaginary cleavage (two each from the two carbon atoms and two from the hydrogen atom). Elemental nitrogen has only three binding electrons. Since electrons are negatively charged, the nitrogen atom would have the charge −3 after the imaginary split. This is therefore also its oxidation number.

All determined oxidation numbers can be added up for checking purposes. Their total must add up to zero if the entire molecule is uncharged.

See also

Individual evidence

  1. 1,01,1 Entry: Oxidation state. In: IUPAC Compendium of Chemical Terminology (the “Gold Book”). doi: 10.1351 / goldbook.O04365 (Version: 2.3.).
  2. ↑ Hans-Dieter Jakubke, Ruth Karcher (Ed.): Chemistry Lexicon, Spektrum Akademischer Verlag, Heidelberg, 2001.
  3. ↑ Entry: Oxidation number. In: IUPAC Compendium of Chemical Terminology (the “Gold Book”). doi: 10.1351 / goldbook.O04363 (Version: 2.3.).
  4. ↑ Karl-Heinz Lautenschläger, Werner Schröter, Joachim Teschner, Hildegard Bibrack, Paperback of Chemistry, 18th edition, Harri Deutsch, Frankfurt (Main), 2001.

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