Oxidation Process

You may thank oxidation if you like oxygen-based cleaners or appreciate hydrogen peroxide’s sterilising properties. On the other hand, oxidation is to blame if you’ve ever dealt with a rusted car or had to throw out browned fruit. Oxidation can occur naturally or be induced. It can be beneficial at times and quite harmful at others.

The loss of electrons is the simplest definition of oxidation. It occurs when one or more electrons are lost by an atom or molecule. Electrons are lost more easily in some elements than in others. It is said that certain elements are easily oxidised. Sodium, magnesium, and iron, for example, are easily oxidised metals.

Elements that are less willing to shed electrons are more difficult to oxidise; they cling to their electrons with a ferocious grip. Nitrogen, oxygen, and chlorine are examples of nonmetals that are difficult to oxidise.

The Process of Oxidation

The characteristics of an atom or molecule alter when it oxidises. When an iron object, for example, gets oxidised, it undergoes transformation as electrons are lost. Oxidized iron is a brittle, reddish powder, while unoxidized iron is a strong, structurally sound metal. What happens to one atom of iron as it oxidises is seen in the diagram below:

Iron takes on a charge after it has been oxidised. Now it has a positive electrical charge of three since three electrons were lost. The number three as well as a positive sign (3+) written as a superscript to the right of the Iron (Fe) symbol are used to signify this positive three charge.

Because iron is easily oxidised, it’s critical to keep it away from oxygen and moisture as much as possible. As long as oxygen is available, iron will continue to lose electrons.

Oxidation-Reduction (Redox) Reactions

The majority of the time, oxidation takes place in conjunction with a process known as reduction. Gaining one or more electrons is the process of reduction. One atom or compound steals electrons from another in an oxidation-reduction reaction, often known as a redox reaction.

Rusting is a well-known redox reaction. Oxygen removes electrons from iron during rusting. The reduction of oxygen occurs while the oxidation of iron occurs. Iron oxide, sometimes known as rust, is the consequence of this reaction. The oxidised type of iron found in rust is distinct from unoxidized, or pure, iron.

For learning the differences between oxidation and reduction, LEO the lion says GER is a highly helpful mnemonic device.

LEO or GER stand for Loss of Electrons (Oxidation) and Gain of Electrons (Reduction) respectively.

Oxidation numbers

Chemists invented the notion of oxidation numbers, which allows them to track electrons before and after a reaction to help identify these less visible redox reactions. The imaginary charge that an atom would have if all of its bonds were entirely ionic is represented by its oxidation number (or oxidation state). Visual scrutiny may be sufficient in some circumstances. For example, we could have deduced that rusting of iron is a redox process simply by seeing the creation of ions from the free components Fe and O2. In other circumstances, however, it is less clear, especially when the reaction only involves nonmetallic components.

Two half-reactions combine to form a full reaction

The oxidation and reduction reaction component of a redox reaction is referred to as a half reaction. The change in oxidation states of the different compounds involved in the redox reaction is used to calculate the half reaction.

For example:

Fe2+ → Fe3+ + e- becomes 2Fe2+ → 2Fe3+ + 2e-

is added to Cl2 + 2e- → 2Cl-

and finally becomes Cl2 + 2Fe2+ → 2Cl- + 2Fe3+

Half-reactions are frequently used to describe what happens in an electrochemical cell like a Galvanic cell battery. Both the metal undergoing oxidation (known as the anode) and the metal experiencing reduction can be described using half-reactions (known as the cathode).

The use of half-reactions to balance redox reactions is common. After balancing the atoms and oxidation numbers in acidic oxidation-reduction processes, the hydrogen ions in the half reaction must be balanced by adding H+ ions.

After balancing the atoms and oxidation numbers, treat the oxidation-reduction processes in basic circumstances as an acidic solution, then add OH- ions to balance the H+ ions in the half reactions (which would give H2O).

Conclusion

Redox reactions, also known as oxidation–reduction processes, are chemical reactions in which electrons are transferred from one species to another. The species that loses electrons is oxidised, whereas the species that gets electrons is reduced. Redox reactions are identified using oxidation numbers, which are assigned to atoms in molecules based on the assumption that all links between them are ionic. During a reaction, oxidation occurs when the oxidation number rises, while reduction occurs when the oxidation number falls.

Remembering with OIL RIG Oxidation and Reduction are two different types of reactions that occur in the body.

So keep in mind that the contemporary definitions of oxidation and reduction are both concerned with electrons (not oxygen or hydrogen). The OIL RIG is a useful tool for remembering which species are oxidised and which are reduced.