Important Topic of Chemistry: Oxidation States

 

The concept of assigning an oxidation state to each atom in a molecule evolved from the chemical bond’s electron-pair concept. The force of attraction exerted by the nuclei of two or more atoms on electrons in the space between them holds atoms within a molecule together. In many cases, electron sharing can be thought of as involving electron-pair bonds between adjacent nuclei. Oxidation States are effective tools in nomenclature, redox chemistry, and other fields. There are a few simple rules for determining the oxidation state of elements in an atom, but they are not universal. In some cases, assigning an oxidation state isn’t useful, or the oxidation state is ambiguous and in particular instances, such as intermetallic compounds, we are better off assigning, rounding off, or estimating oxidation states ourselves.

Overview

Overall, oxidation state or number aids in the description of electron transfer. The oxidation number/state is also used to calculate the changes in redox reactions. Meanwhile, it’s very similar to valence electrons. It is prevalent to regard a single value of electronegativity as valid for the majority of bonding situations that a given atom can be in. While this approach has the advantage of being simple, it is clear that an element’s electronegativity is not an invariable atomic property; rather, it can be thought of as depending on a quantity known as the element’s oxidation number.

The assignment of oxidation numbers is one method for characterising atoms in a molecule and keeping track of electrons. The oxidation number is the electric charge that an atom would have if the bonding electrons were exclusively assigned to the more electronegative atom, and it can be used to determine which atom is oxidised and which is reduced in a chemical process.

Oxidation states

The oxidation state, also known as the oxidation number is an atom’s hypothetical charge if all of its bonds to other atoms were fully ionic. It expresses the degree of oxidation (electronic loss) of an atom in a chemical compound. In fact, the oxidation state can be positive, negative, or zero. While fully ionic bonds do not exist in nature, many bonds have high ionicity, making oxidation state a useful predictor of charge.

An atom’s oxidation state does not represent its “real” formal charge or any other actual atomic property. This is especially true in high oxidation states, where the ionisation energy required to generate a multiply positive ion is far greater than the energies available in chemical reactions. Furthermore, the oxidation states of atoms in a compound may differ depending on the electronegativity scale used in their calculation. As a result, an atom’s oxidation state in a compound is purely a formality. It is, however, necessary for understanding the nomenclature conventions of inorganic compounds. In addition, several observations about chemical reactions can be explained at a basic level using oxidation states.

In most cases, oxidation states are represented by integers that can be positive, zero, or negative. In some cases, an element’s average oxidation state is a fraction.

The oxidation state is represented in inorganic nomenclature by a Roman numeral placed after the element name inside the parenthesis or as a superscript after the element symbol, for example, Iron(III) oxide.

The term “oxidation” was formulated by Antoine Lavoisier to describe the reaction of a substance with oxygen. Much later, it was discovered that the substance loses electrons when oxidised, and the meaning was expanded to include other reactions in which electrons are lost, regardless of whether oxygen is present. The process of increasing an atom’s oxidation state through a chemical reaction is known as oxidation; the process of decreasing an atom’s oxidation state is known as reduction. Of this kind reactions involve the formal transfer of electrons: a net gain in electrons is referred to as reduction, while a net loss of electrons is referred to as oxidation. The oxidation state of pure elements is zero.

The oxidation state of an atom is not considered to be the atom’s true charge. When there is an increase in oxidation state in a chemical reaction, this is referred to as oxidation; when there is a decrease in oxidation state, this is referred to as reduction. For carbon in CH4, the lowest known oxidation state is 4. (methane). In tetroxoiridium, the highest known oxidation state is +9. (IX).

To calculate the oxidation number, we must first understand and then apply certain rules. There are six ground rules:

1. Every atom in an element, whether free or combined, has an oxidation number of zero. Each atom in H₂, Cl₂, P₄, Na, Al, O₂, O₃, S₈, and Mg clearly has an oxidation number of zero.
2. The oxidation number of ions with only one atom is equal to the ion’s actual charge.
3. The oxidation number of oxygen in the majority of compounds is –2. This rule has two exceptions.

Peroxides- Each oxygen atom has an oxidation number of –1. For example, in the case of Na₂O₂

Superoxides: each oxygen atom is assigned an oxidation number of –(1/2).

For example, KO2 and dioxygen difluoride (oxygen is bonded to fluorine), where the oxygen atom has an oxidation number of +1.

4. Except when bonded to metals containing two elements, hydrogen’s oxidation number is +1. CaH2, for example, has an oxidation number of –1.
5. When fluorine and other halogens appear as halide ions in their compounds, they have an oxidation number of –1. The oxidation number of iodine, chlorine, and bromine, when combined with oxygen, is positive.
6. Whenever the oxidation numbers of a compound’s atoms are added together, the algebraic sum must equal zero. When the oxidation numbers of an ion’s atoms are added together in the case of a polyatomic ion, the algebraic sum must equal the ion’s charge. Consider (CO₃)^2–: the algebraic sum of the oxidation numbers of one carbon atom and three oxygen atoms is -2.

Oxidation number of sulphur

The sulphur atom in the SO₄^2- ion must have an oxidation number of +6. Sulphur has a large number of oxidation numbers depending on the compound formed.

As an example: In H_₂SO_₄, the oxidation number of S is +6.

Let X denote the number of oxidations of S in H_₂SO_₄.

The oxidation number of hydrogen is +1, while that of oxygen is -2.

2(+1)+X+4(−2)=0

2+X−8=0

X−6=0

X=+6

Oxidation number of oxygen

The oxidation number of oxygen in the majority of compounds is –2. This rule has two exceptions.

Peroxides- Each oxygen atom has an oxidation number of –1. For example, in the case of Na2O2

Superoxides: each oxygen atom is assigned an oxidation number of –(1/2).

For example, KO₂ and dioxygen difluoride (oxygen is bonded to fluorine), where the oxygen atom has an oxidation number of +1.

FAQ’s

What is meant by oxidation number?

The total number of electrons that an atom gains or loses in order to form a chemical bond with another atom is referred to as the oxidation number, also known as the oxidation state.

What is a positive oxidation state?

When the oxidation number is positive, the atom loses electrons; when it is negative, the atom gains electrons. Calcium has a charge of +2, indicating that two electrons have been lost. Oxygen's -2 charge indicates that it has gained two electrons.

Does fluorine show a positive oxidation state?

Fluorine is said to be the most electronegative element in the periodic table. As a result, fluorine is always in a negative oxidation state and can never be in a positive one.