Reducing agent

A reducing agent (also called a reductant or reducer) is the element or compound in a reduction-oxidation (redox) reaction that donates an electron to another species; however, since the reducer loses an electron we say it is "oxidized". This means that there must be an "oxidizer"; because if any chemical is an electron donor (reducer), another must be an electron recipient (oxidizer). Thus reducers are "oxidized" and oxidizers are "reduced". For example, consider the following reaction:


 * 2 +  → 2  + 2

The reducing agent in this reaction is ferrocyanide. It donates an electron, becoming oxidized to ferricyanide, simultaneously the oxidizer chlorine is reduced to chloride.

In organic chemistry, reduction more specifically refers to the addition of hydrogen to a molecule, though the aforementioned definition still applies. For example, benzene is reduced to cyclohexane in the presence of a platinum catalyst:
 * C6H6 + 3 H2 → C6H12

In organic chemistry, good reducing agents are reagents that deliver H2.

Characteristics of reducing agents
Strong reducing agents easily lose (or donate) electrons. An atom with a relatively large atomic radius tends to be a better reductant. In such species, the distance from the nucleus to the valence electrons is so long that these electrons are not strongly attracted. These elements tend to be strong reducing agents. Good reducing agents tend to consist of atoms with a low electronegativity, the ability of an atom or molecule to attract bonding electrons, and species with relatively small ionization energies serve as good reducing agents too. "The measure of a material to oxidize or lose electrons is known as its oxidation potential". The table below shows a few reduction potentials that could easily be changed to oxidation potential by simply reversing the sign. Reducing agents can be ranked by increasing strength by ranking their oxidation potentials. The reducing agent is stronger when it has a more positive oxidation potential and weaker when it has a negative oxidation potential. The following table provides the reduction potentials of the indicated reducing agent at 25 °C.

To tell which is the strongest reducing agent, one can change the sign of its respective reduction potential to make it oxidation potential. The bigger the number, the stronger the reducing agent. For example, among Na, Cr, Cu and Cl−, Na is the strongest reducing agent and Cl− is the weakest one.

Common reducing agents include metals potassium, calcium, barium, sodium and magnesium, and also compounds that contain the H− ion, those being NaH, LiH, LiAlH4 and CaH2.

Some elements and compounds can be both reducing or oxidizing agents. Hydrogen gas is a reducing agent when it reacts with non-metals and an oxidizing agent when it reacts with metals.


 * 2 Li(s) + H2(g) → 2 LiH(s)

Hydrogen acts as an oxidizing agent because it accepts an electron donation from lithium, which causes Li to be oxidized.

Half reactions: 2 Li(s)0 → 2 Li(s)+ + 2 e−::::: H20(g) + 2 e− → 2 H−(g)


 * H2(g) + F2(g) → 2 HF(g)

Hydrogen acts as a reducing agent because it donates its electrons to fluorine, which allows fluorine to be reduced.

Half reactions: H20(g) → 2 H+(g) + 2 e−::::: F20(g) + 2 e− → 2 F−(g)

Importance of reducing and oxidizing agents
Reducing agents and oxidizing agents are the ones responsible for corrosion, which is the “degradation of metals as a result of electrochemical activity”. Corrosion requires an anode and cathode to take place. The anode is an element that loses electrons (reducing agent), thus oxidation always occurs in the anode, and the cathode is an element that gains electrons (oxidizing agent), thus reduction always occurs in the cathode. Corrosion occurs whenever there’s a difference in oxidation potential. When this is present, the anode metal begins deteriorating, given there is an electrical connection and the presence of an electrolyte.

Example of redox reaction
The formation of iron(III) oxide;
 * 4Fe + 3O2 → 2Fe2O3

In the above equation, the Iron (Fe) has an oxidation number of 0 before and 3+ after the reaction. For oxygen (O) the oxidation number began as 0 and decreased to 2−. These changes can be viewed as two "half-reactions" that occur concurrently:


 * 1) Oxidation half reaction: Fe0 → Fe3+ + 3e−
 * 2) Reduction half reaction: O2 + 4e− → 2 O2−

Iron (Fe) has been oxidized because the oxidation number increased and is the reducing agent because it gave electrons to the oxygen (O2). Oxygen (O2) has been reduced because the oxidation number has decreased and is the oxidizing agent because it took electrons from iron (Fe)

Common reducing agents

 * Lithium aluminium hydride (LiAlH4)
 * Nascent (atomic) hydrogen
 * Sodium amalgam
 * Sodium borohydride (NaBH4)
 * Compounds containing the Sn2+ ion, such as tin(II) chloride
 * Sulfite compounds
 * Hydrazine (Wolff-Kishner reduction)
 * Zinc-mercury amalgam (Zn(Hg)) (Clemmensen reduction)
 * Diisobutylaluminum hydride (DIBAH)
 * Lindlar catalyst
 * Oxalic acid (C2H2O4)
 * Formic acid (HCOOH)
 * Ascorbic acid (C6H8O6)
 * Phosphites, hypophosphites, and phosphorous acid
 * Dithiothreitol (DTT) – used in biochemistry labs to avoid S-S bonds
 * Compounds containing the Fe2+ ion, such as iron(II) sulfate