Each atom has an oxidation number during the oxidation-reduction reaction, which shows its ability to gain or donate electrons. E.g., the Fe³⁺( iron ion) possesses an oxidation number of +3 as it can gain three electrons to form a chemical bond. In comparison, the O²⁻ (oxide ion) has an oxidation number of −2 because it has accepted two electrons. In a neutral substance, the sum of the oxidation numbers is zero; e.g., in hematite (Fe₂O₃), the oxidation state of the both the iron atoms is +3 and oxidation number of the all the three oxygen atom is -2 . 

Discovery and Representation of oxidation number

Antoine Lavoisier created the oxygen-based dualism theory of chemistry, which gave rise to the concept of oxidation states. At this period, the terms oxidation and reduction first arise in the literal sense of an element reacting with oxygen and vice versa. Despite the findings of several scientists, no attempt was made to stop using the terms oxidation and reduction to describe reactions involving salts and other compounds that now contain no oxygen. 

After discovering the new ionic theory of dissociation in electrochemistry and the electronic theory of bonding and structure, chemists decided that the growth of positive valency of an atom corresponds to oxidation and the development of negative valency corresponds to reduction. After that, scientists began referring to an element’s multiple oxidation states and officially coined oxidation number or oxidation state and the parallel term oxidation potential.

According to inorganic nomenclature, the oxidation state is represented by a Roman numeral placed after the element name inside the parenthesis, e.g., Iron(III) oxide.

What is Oxidation?

Oxidation is the process in which the movement of electrons is included. When a substance donates electrons, it is called to be oxidised. E.g., oxidation of magnesium occurs when magnesium metal and oxygen react and form magnesium oxide.

The oxidation state is zero in the case of an uncombined element. The element has not been oxidised yet, e.g., Xe or Cl2 or S8. In a neutral compound, the sum of the oxidation states of all the atoms is zero. In a charged compound, the sum of oxidation states of all the atoms is equal to the charge on the ion. A substance’s more electronegative element is given a negative oxidation state. A positive oxidation state is assigned to the less electronegative one. The most electronegative element is fluorine, followed by oxygen.

Using oxidation states as a tool when it comes to naming substances

Iron(II) sulphate and iron(III) chloride are two examples of compounds you may have encountered. The oxidation states of iron in the two compounds are (II) and (III), respectively: +2 and +3. This indicates that Fe²⁺ and Fe³⁺ ions are present.

This can be applied to the negative ion as well. FeSO₄ is the chemical formula for iron(II) sulphate. FeSO₄, also known as iron(II) sulfite, is another chemical. The contemporary names refer to the two compounds’ oxidation states of sulphur.

SO₄^2- is the sulphate ion. The oxidation number of sulphur is +6. The ion is known as the sulphate (VI) ion.

SO₃^2- is the sulfite ion. The oxidation state of sulphur is +4. The sulphate (IV) ion is the proper name for this ion. The -ate suffix simply indicates that the sulphur is in a negative ion state.

As a result, FeSO₄ is iron(II) sulphate(VI), while FeSO₃ is iron(II) sulphate (IV). 

Highest and lowest oxidation states

In the case of ruthenium, xenon, osmium, hassium, iridium, and some complexes involving plutonium, the highest oxidation state is +8. For some elements from the carbon group, the oxidation state is -4. Plutonium changes colour according to oxidation state.

How to calculate Oxidation Number?

• Each atom in a free or uncombined element has an oxidation number of zero. In P₄, H₂, Cl₂, Na, Al, O₂, O₃, S, and Mg, each atom has an oxidation number of zero
• The actual charge of an ion is equal to the oxidation number of ions with only one atom
• The oxidation number of oxygen in most substances is –2. In this case, there are two exceptions. Peroxides have an oxidation number of –1 assigned to each oxygen atom. Superoxides, for example, have an oxidation number of –(1/2) for each oxygen atom. KO₂ oxygen is coupled to fluorine in dioxygen difluoride, for example, where the oxygen atom has an oxidation number of +1.
• The hydrogen oxidation number is +1 unless it is linked to two-element metals. CaH₂, for example, has an oxidation number of –1
• When fluorine and other halogens exist as halide ions in their compounds, they have an oxidation number of –1. The oxidation number of iodine, chlorine, and bromine, when mixed with oxygen, is positive
• When the oxidation numbers of the atoms of a compound are put together, the algebraic total must equal zero. When the oxidation numbers of an ion’s atoms are put together in the case of polyatomic ions, the algebraic total must match the ion’s charge. Consider CO₃^2-: the algebraic sum of one carbon atom and three oxygen atom’s oxidation numbers is -2

Assigning Oxidation States Rules

The amount of electrons, e-, that an atom loses, obtains or seems to utilise as it unites with other atoms in a molecule determines its oxidation status (OS). There are seven rules to follow when identifying an atom’s OS:

1. An individual atom’s OS is 0.
2. In a neutral species, the total OS of all atoms is 0; in an ion, the total OS is equal to the charge on ion.
3. Metals in Group 1 have an OS of +1, while metals in Group 2 have an OS of +2.
4. In compounds, fluorine’s OS is -1.
5. In most compounds, hydrogen has an OS of +1.
6. In most compounds, oxygen has an OS of -2.
7. Group 17 elements have an OS of -1 in binary metal compounds, Group 16 elements have an OS of -2, and Group 15 elements have an OS of -3.

Conclusion

The total number of electrons that have been removed from an element (forming a positive oxidation state) or added to an element (generating a negative oxidation state) to get it to its current state is the oxidation state of an atom. An increase in oxidation state is referred to as oxidation. Reduction is the process of lowering the oxidation state of a substance. Understanding the concept of oxidation states begins with recognising this simple pattern. Without using electron-half-equations, the change in the oxidation state of an element during a process determines whether it has been oxidised or reduced.