Oxidation State

The chemical process of oxidation-reduction, better known as redox reaction is a common phenomenon around the world. It is also an essential contributor to the metabolic process, whereby nutrient oxidation leads to energy release and enables life forms to thrive. Exposure of various elements, as well as compounds, causes combustion and release of water, carbon dioxide, and energy. Thus, to gain a better understanding of reactions like redox and combustion, one must be aware of oxidation state or OS, a chemical characteristic exhibited by various elements.

The term redox is created by combining two words RED for reduction and OX from oxidation. Where two major processes of reduction and oxidation took place simultaneously and satisfy the molecular requirement of each other.

The displacement reactions are classic examples of redox reactions where one species is oxidized and loses electrons while the other is being reduced by gaining that same electron.

What does Oxidation State Mean?

The oxidation number or the oxidation state is theoretically a charge of an atom if all of its bonds to several different atoms were fully ionic. They define the level of oxidation of an atom in a chemical compound.

On the conceptual level, the oxidation state can be either represented by using integrals like positive, negative, or zero.

In other words, OS refers to a specific number assigned to elements in different chemical combinations. These numbers are a representation of electron quantity lost or gained by an element’s atom to result in a chemical bond with another element. Also interchangeably used with oxidation number, it is used for the determination of changes that are taking place in a redox reaction. It has a similar numerical representation as valence electrons but is typically differentiated from formal charge. To better understand the oxidation state definition, you must also know that elements must act as a reducing agent on oxidation, resulting in the release of electrons.

Origin and Discovery of Oxidation State

The popular French chemist, Antoine Lavoisier first used oxidation as a means to describe the reaction of oxygen with any substance. Later experiments remarkably identified oxidation to result in electron loss. The use of the term oxidation was thus expanded to other reactions that identified electron loss, whether or not it involved oxygen, et. Thus, it increased its scope of usage. Such electron losses were thus identified with values known as oxidation state. One could thus define oxidation number or state by putting a value to such electron losses during a reaction, which usually stood as integers. At times, the OS can also be represented as a fraction. For instance, the OS of iron in \[Fe_{3}O_{4}\] is valued at 8/3. Before moving on to more about oxidation number or state, take a brief look at the process of oxidation.

In simple words, Lavoisier uses the term oxidation for the first time in history. It represents the reaction of a substance with oxygen.

After a long period of time, it was discovered that the substance, upon being oxidized, loses its electrons. It further extends to include other reactions where electrons are lost. Except for the fact whether oxygen was involved or not.

What does Oxidation Entail?

The contact between the oxygen molecules and substances causes an Oxidation reaction. It can be simply described as an atom's enhanced oxidation state through a chemical reaction. It’s the exact contrary part of the reduction reaction. In both reactions, the transfer of electrons is required.

Oxidation refers to a chemical reaction that involves electron movement between the elements of any compound. The process’s character is exhibited when an element donates electrons. It is also denoted by an increased oxidation state. A common example of oxidation is the reaction of iron (Fe) with oxygen (\[O_{2}\]). The reaction of these two elements results in the formation of rust, whereby the electrons lost by iron are gained by oxygen.

Oxidation States – The Highest and the Lowest

While oxidation involves the increase in oxidation state, its decrease is denoted by reduction. However, there have been quantity limits identified for the state, whereby the highest OS has been marked at +9 for tetrox iridium and the lowest in the case of carbon in methane or \[CH_{4}\] at -4.

Identifying oxidation states in elements or compounds is based on the following few rules one must take care of. You must also follow these rules in the right order and consider the one appearing first in order in the case of conflict.

Rules to identify and assign Oxidation States

Some other examples that satisfy the OS rules are mentioned below.

1. In \[NH{_{4}}^{+}\], the oxidation state of nitrogen is -3 and that of hydrogen is -1.

2. In \[SO{_{4}}^{2-}\], oxygen’s OS is -2 and sulphur, +6.

3. For \[OH^{-}\], OS for oxygen stands at -2 and that for hydrogen at -1.

Also, while the oxidation states are mostly represented by integers, some can also have fractional values. Nevertheless, the rules still apply when assigning the OS to elements in a given compound. If the ions of an element in a given compound can be separated, the OS is determined by the process of fragmenting, which can result in fractional values for the element.

Quick Question 1: Does oxidation always involve oxygen?

Ans: Oxidation happens when an ion atom loses electrons (one or more) in the chemical reaction. During the oxidation process, the oxidation state of the chemical species increases. This reaction doesn’t require oxygen necessarily.

Basically, the term is used when oxygen happens to lose an electron during the reaction.

Oxidation Number Calculation

You can calculate an atom’s oxidation number in a given compound by taking care of the following few computation rules.

1. Remember rule 1 suggests every atom that is free or non-combined carry an OS of ‘0’. As per rule 2, ions with only one atom carry a charge that equals the ion’s OS.

• While the OS in the case of oxygen stands at -2, some exceptions to this rule include – 

• Each oxygen atom carries an OS of -1 in the case of peroxides.

2. Oxygen atoms carry an OS of –(1/2) in the case of superoxides.

3. When forming a bond with fluorine, the OS of oxygen can vary.

4. For hydrogen bonding with metals carrying two elements, OS is -1.

5. OS is positive for halogens like chlorine, iodine, and bromine when combined with oxygen.

6. The sum of oxidation numbers added for different atoms in a compound should equal zero.

7. The OS of an ionic compound equals the ion’s charge when calculating it for polyatomic ions.

Quick Practice Questions

Assign oxidation numbers to the elements of the following compounds using the rules and the calculation method given above.

• \[H_{2}SO_{4}\]

• \[CuSO_{4}\]

• \[FeCl_{3}\]

• \[S_{2}O{_{3}}^{2-}\]

• NaH

• The difference between the oxidation number and oxidation state are terms with interchangeable use, they are not exactly the same. The following points state the difference between oxidation state and oxidation number.

• Although oxidation number and oxidation state are terms with interchangeable use, they are not exactly the same. The following points state the difference between oxidation state and oxidation number.

Oxidation States and Numbers:

Now, take a look at some other essential details on oxidation state and numbers that can help in the calculation.

When expressed in terms of the periodic table, the oxidation numbers of various elements can be identified based on the following properties.

• Alkali metals of Group 1A carry an OS of -1.

• Alkali earth metals of Group 2A carry an OS of -2.

• Elements in Group 3A often exhibit an OS of +3. They can also have an OS of +2 or +1.

• OS for elements in Group 3A remains in the range of -4 and +4. (Image to be added soon)

• In the case of non-metals, the OS can vary on a wide range. However, when calculated, subtracting the minimum OS from the maximum results in +8.

• The maximum OS of an element can go up to 7 subtracted by its group number.

• Elements of Group 1 have an OS of +1.

• For Group 2, it is +2 and for Group 3, +3.

• Group 4 elements range between -4 and +4 while Group 5 elements are between -3 and +5. Group 6 and 7 elements range from -2 to +6 and -1 to +7.

• For inert gases, the OS is always zero.

While this was all about oxidation state, its exhibition, and properties for different elements, you can carry on with your exam preparation by referring to chemistry solutions as well. Vedantu hosts a range of study material online to assist students with thorough preparation. 

Download our Vedantu app today for a streamlined approach to exam preparation.

Quick Question 2: Does oxidation produce heat?

Ans: The Oxidation reaction is a reaction where there is a loss of electrons, generally through the addition of oxygen or reduction of hydrogen. The oxidation of glucose to pyruvate emits energy, which would be listed as heat, but there is no solid proof for this. 

Types of Redox Reactions

There are five main types of redox reactions mentioned as follows:

• Combination Reactions

A combination reaction involves the combination of two different compounds. By this combination, a singular compound is formed. 

\[ X + Y \rightarrow  XY\]

Here are some of the common examples of combination reactions.

\[4Fe + 3O_{2} \rightarrow  2Fe_{2}O_{3}\]

\[2Ca + O_{2} \rightarrow  2CaO\]

• Decomposition Reactions

In this type of redox reaction, the breakdown of the compound happens into various compounds. The decomposition reactions are the contrary of combination reactions.

Here are some basic examples of these types of reactions:

\[2H_{2}O \rightarrow  2H_{2} + O_{2}\]

\[H_{2}CO_{3} \rightarrow  H_{2}O + CO_{2}\]

In the above reaction, the total compound is broken down into the small chemical compound in the form of XY → X + Y

In some specific cases, it’s deduced that the decomposition reactions do not fulfill the criteria of redox reactions.

One such example is as follows:

\[CaCO_{3} \rightarrow  CaO + CO_{2}\]

• Displacement Reactions

In this type of reaction, an ion, atom, or molecule in a compound is interchanged by an ion atom or molecule in another compound. Simply, it is an exchange deal between compounds. 

For example \[A + BC \rightarrow  AC + B\]

It can also be categorized into the following sub-category. 

1. Metal displacement Reaction

2. Non-metal displacement Reaction

1. Metal Displacement

In this kind of displacement reaction, one metal present in a compound is displaced by the metal present in another metal. These kinds of reactions are basically used in the metallurgical processes where pure metals are extracted from their ores.

Here’s a common example of a metal displacement reaction

\[CuSO_{4} + Zn \rightarrow  Cu + ZnSO_{4}\]

2. Non-Metal Displacement

Non-Metal Displacement reactions involved hydrogen displacement. There are rare cases where oxygen displacement happens in these types of reactions. 

• Combustion Reactions

This type of redox reaction happens between an oxidant and a fuel. By using atmospheric oxygen. This produces a gaseous substance called smoke.

Here is an example of a reaction.

\[CH_{4}(g) + 2O_{2}(g) \rightarrow  CO_{2}(g) + 2 H_{2}O(g)\]

• Disproportionation Reactions

In the disproportionation reactions, a single reactant is oxidized and reduced. It’s also called a Dismutation reaction. 

Here’s a convenient example:

\[Mn_{2}O_{3} \rightarrow  Mn_{2} + MnO_{2}.\]

In this reaction, one compound of intermediate oxidation state breaks into two compounds, where there is a difference between oxidation states i.e. one compound is of higher oxidation states while the other is lower oxidation states.