Shows iodine an oxidation of +1
The Oxidation numberNox (also oxidation state, oxidation value) indicates how many elementary charges an atom has formally taken up or given off within a compound, for example in a redox reaction. It therefore corresponds to the hypothetical ion charge of an atom in a molecule or the actual charge of monatomic ions.
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.
The oxidation number is a useful formalism for chemical considerations that often has little to do with the real charge of an atom. It is quite possible that atoms in a compound are assigned a negative formal oxidation number, although they also carry a positive formal charge. In covalent compounds, the oxidation number differs from the concept of valence.
The oxidation numbers are used in redox reactions to better recognize the processes. The transfer of electrons from one atom to another is shown by the fact that the oxidation number of one (which emits electrons) increases and that of the other (which accepts electrons) decreases. Often it is only through the determination of the oxidation numbers of individual atoms that it becomes clear which chemical reaction is taking place.
Specification of the oxidation number
Oxidation numbers are written in Roman numerals over the atomic symbols in compounds (e.g. O−II). If the element symbol is on its own, they are often written as Arabic numerals as in the case of ions. According to IUPAC, signs are only set for negative oxidation numbers.
Determination of the oxidation number
The oxidation number can be derived using the following rules:
- Atoms in the elementary state always have the oxidation number 0 (0 is also possible in compounds).
- In the case of monatomic ions, the oxidation number corresponds to the ion charge.
- The sum of the oxidation numbers of all atoms of a polyatomic neutral compound is equal to 0.
- The sum of the oxidation numbers of all atoms of a polyatomic ion is equal to the total charge of this ion.
- 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.
- Most elements can occur in several oxidation states.
In practice, it has proven helpful to formulate a few rules for determining the oxidation number:
- The fluorine atom (F) as an element with the highest electronegativity always has the oxidation number −I in compounds.
- Oxygen atoms get the oxidation number −II - except in peroxides (then: −I) and in connection with fluorine (then: + II).
- Other halogen atoms (such as chlorine, bromine, iodine) generally have the oxidation number (−I), except in connection with oxygen or a halogen that is higher in the periodic table.
- In compounds, metal atoms always have a positive oxidation number as ions.
- Alkali metals always have + I and alkaline earth metals always + II as the oxidation number.
- Hydrogen atoms get the oxidation number + I, except when hydrogen is directly linked to more “electropositive” atoms such as metals (hydrides) or to itself).
- In the elementary state, the oxidation number is always 0 (e.g. I.2, C, O2, P4, S.8).
- In ionic compounds (salts) the sum of the oxidation numbers is identical to the ionic charge.
- In covalent connections (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.
- The highest possible oxidation number of an element corresponds to the number of major or minor groups in the periodic table (PSE).
Graphic determination of oxidation numbers
As an example, the phosphoric acid (H.3PO4) serve:
- First the Lewis formula is recorded.
- Then the electrons are assigned to the atoms according to electronegativity
- The oxidation number can then be calculated based on the valence electrons. Example: Oxygen normally has 6 valence electrons (VI. Main group). Due to the higher electronegativity of oxygen, the binding electrons between oxygen and hydrogen (or phosphorus) can be assigned to oxygen. In the balance sheet, the oxygen receives two additional electrons in addition to the 6 available. Hence the oxidation number is −II. Phosphorus is in main group V, so it normally has 5 valence electrons. Since these are all assigned to oxygen, it “lacks” five electrons and it has the oxidation number + V.
Another example shows on the one hand how one and the same atom (here the carbon atom) has different oxidation numbers, and on the other hand how oxidation numbers change during the reaction (here using the example of the Tollens sample / silver mirror sample):
Categories: Chemical Bond | Chemical reaction
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