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In chemistry, a reducing agent (also known as a reductant, reducer, or electron donor) is a chemical species that "donates" an electron to an electron recipient (called the oxidizing agent, oxidant, oxidizer, or electron acceptor). 
Examples of substances that are common reducing agents include the alkali metals, formic acid, oxalic acid, and sulfite compounds. 
In their pre-reaction states, reducers have extra electrons (that is, they are by themselves reduced) and oxidizers lack electrons (that is, they are by themselves oxidized). This is commonly expressed in terms of their oxidation states. An agent's oxidation state describes its degree of loss of electrons, where the higher the oxidation state then the fewer electrons it has. So initially, prior to the reaction, a reducing agent is typically in one of its lower possible oxidation states; its oxidation state increases during the reaction while that of the oxidizer decreases. 
Thus in a redox reaction, the agent whose oxidation state increases, that "loses/donates electrons", that "is oxidized", and that "reduces" is called the reducer or reducing agent, while the agent whose oxidation state decreases, that "gains/accepts/receives electrons", that "is reduced", and that "oxidizes" is called the oxidizer or oxidizing agent. 
For example, consider the overall reaction for aerobic cellular respiration:Importance
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.Examples of redox reaction
Historically, reduction referred to the removal of oxygen from a compound, hence the name 'reduction'. An example of this phenomenon occurred during the Great Oxidation Event, in which biologically−produced molecular oxygen (dioxygen (O2), an oxidizer and electron recipient) was added to the early Earth's atmosphere, which was originally a weakly reducing atmosphere containing reducing gases like methane (CH4) and carbon monoxide (CO) (along with other electron donors) and practically no oxygen because any that was produced would react with these or other reducers (particularly with iron dissolved in sea water), resulting in their removal. 
By using water as a reducing agent, aquatic photosynthesizing cyanobacteria produced this molecular oxygen as a waste product. This O2 initially oxidized the ocean's dissolved ferrous iron (Fe(II) − meaning iron in its +2 oxidation state) to form insoluble ferric iron oxides such as Iron(III) oxide (Fe(II) lost an electron to the oxidizer and became Fe(III) − meaning iron in its +3 oxidation state) that precipitated down to the ocean floor to form banded iron formations, thereby removing the oxygen (and the iron). The rate of production of oxygen eventually exceeded the availability of reducing materials that removed oxygen, which ultimately led Earth to gain a strongly oxidizing atmosphere containing abundant oxygen (like the modern atmosphere). The modern sense of donating electrons is a generalization of this idea, acknowledging that other components can play a similar chemical role to oxygen.
The formation of iron(III) oxide;Oxidation half reaction: Fe0 → Fe3+ + 3e−
Reduction half reaction: O2 + 4e− → 2 O2−Iron (Fe) has been oxidized because the oxidation number increased. Iron 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), a very strong reducing agent
Red-Al (NaAlH2(OCH2CH2OCH3)2), a safer and more stable alternative to lithium aluminum hydride
Hydrogen without or with a suitable catalyst; e.g. a Lindlar catalyst
Sodium amalgam (Na(Hg))
Sodium-lead alloy (Na + Pb)
Zinc amalgam (Zn(Hg)) (reagent for Clemmensen reduction)
Diborane
Sodium borohydride (NaBH4)
Ferrous compounds that contain the Fe2+ ion, such as iron(II) sulfate
Stannous compounds that contain the Sn2+ ion, such as tin(II) chloride
Sulfur dioxide (sometimes also used as an oxidizing agent), Sulfite compounds
Dithionates, e.g. Na2S2O6
Thiosulfates, e.g. Na2S2O3 (mainly in analytical chemistry)
Iodides, such as potassium iodide (KI) (mainly in analytical chemistry)
Hydrogen peroxide (H2O2) – mostly an oxidant but can occasionally act as a reducing agent (typically in analytical chemistry.)
Hydrazine (Wolff-Kishner reduction)
Diisobutylaluminium hydride (DIBAL-H)
Oxalic acid (C2H2O4)
Formic acid (HCOOH)
Ascorbic acid  (C6H8O6)
Reducing sugars, such as erythrose, see Aldose
Phosphites, hypophosphites, and phosphorous acid
Dithiothreitol (DTT) – used in biochemistry labs to avoid SS-bonds
Carbon monoxide (CO)
Cyanides in hydrochemical metallurgical processes
Carbon (C)
Tris-2-carboxyethylphosphine hydrochloride (TCEP)
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