Reducing agents

In a redox reaction, a reducing agent (electron donor, reductant, or reducer) is a molecule or element that donates or loses an electron to an electron recipient (also known as the oxidizing agent, oxidizer, or oxidant). The oxidation state of the reducer grows while the oxidation state of the oxidizer drops. This is stated by saying that reductants experience oxidation and get oxidized, whilst oxidants experience reduction and get reduced. Thus, reducers reduce oxidants by lowering their oxidation state, whereas oxidizers oxidize reductants by raising their oxidation state.
Reducers have additional electrons in their before reaction states, whereas oxidants lack electrons . A reducing substance is known as an electron donor when it is in one of its lowest possible oxidation states. Earth metals, oxalic acid, formic acid, and sulfite compounds are examples of reducing agents. Consider the whole response to cellular respiration i.e. aerobic in nature:

C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l)

Because the oxygen is being reduced, it serves as the oxidizing agent. The reducing agent is glucose which is being oxidized.
Though the aforementioned concept still applies, inorganic chemistry, reduction generally refers to the addition of hydrogen to a molecule. For example, in the presence of a platinum catalyst, the oxidizing agent benzene is reduced to cyclohexane:

C6H6 + 3 H2 → C6H12

Features of reducing agents

Take into account the following reaction

2 [Fe(CN)6]4− + Cl2 → 2 [Fe(CN)6]3− + 2 Cl−

• Ferrocyanide ([Fe(CN)6]4) serves as the reducing agent in this process.
• It provides an electron and undergoes oxidation to ferricyanide ([Fe(CN)6]3).
• The electron is simultaneously accepted by the oxidizer chlorine (Cl 2), which is reduced to chloride (Cl).
• lectrons are quickly lost by strong reducing agents.
• ?An atom with a high atomic radius is a superior reductant.
• The distance between the nucleus and the valence electrons in such species is so great that these electrons are not significantly attracted.
• These elements are known to be powerful reducing agents.
• Good reducing agents are composed of atoms with low electronegativity, which is the capacity of a molecule or an atom to attract bonding electrons, as well as species with relatively low ionisation energies.
• The reduction potential of a substance is a measure of its capacity to decrease.
• When the reducing agent has a higher negative reduction potential, it is stronger; when it has a higher positive reduction potential, it is weaker.
• The stronger the species' affinity for electrons and inclination to be reduced, the higher the reduction potential.

Metals such as potassium, calcium, magnesium. barium, and sodium, are common reducing agents. The compounds containing the H ion, such as LiH, NaH, LiAlH4, and CaH2 also fall in the same category.

Some elements and compounds have the ability to act as both reducing and oxidising agents. When hydrogen gas reacts with nonmetals, it is a reducing agent, and when it reacts with metals, it is an oxidising agent.

2 Li(s) + H2(g) → 2 LiH(s)

Because it receives an electron donation from the reducing agent lithium (whose reduction potential is -3.04), hydrogen functions as an oxidising agent, causing Li to be oxidised and Hydrogen to be reduced.

H2(g) + F2(g) → 2 HF(g)

Because it provides electrons to fluorine, hydrogen works as a reducing agent, allowing fluorine to be reduced.

Significance

Corrosion is caused by reducing and oxidising substances, which are responsible for the degradation of metals as a result of electrochemical activity. Corrosion requires an anode and a cathode to take place. The anode is a reducing agent that loses electrons, so oxidation always happens in the anode. The cathode is an oxidising agent that receives electrons, so reduction always occurs in the cathode. Corrosion happens wherever there is an oxidation potential difference. Given an electrical connection and the presence of an electrolyte, the anode metal begins to deteriorate when this happens.

< The synthesis of iron(III) oxide

4Fe + 3O2 → 4Fe3+ + 6O2− → 2Fe2O3

Iron has an oxidation number of 0 before and 3+ after the reaction in the preceding equation. The oxidation number for oxygen began at 0 and dropped to 2. These modifications may be thought of as two half-reactions that happen at the same time:
1. Oxidation half reaction: Fe0 → Fe3+ + 3e−
2. Reduction half reaction: O2 + 4e− → 2 O2−

Because the oxidation number increased, iron (Fe) was oxidised. Because it provided electrons to oxygen, iron is the reducing agent (O2). Because the oxidation number has lowered, oxygen (O2) has been reduced and is the oxidising agent because it has taken electrons from iron (Fe).