Reducing agent has effect reduction
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:2 [Fe(CN)6]4− + Cl2 → 2 [Fe(CN)6]3− + 2 Cl−The reducing agent in this reaction is ferrocyanide ([Fe(CN)6]4−). It donates an electron, becoming oxidized to ferricyanide ([Fe(CN)6]3−). Simultaneously, that electron is received by the oxidizer chlorine (Cl2), which is reduced to chloride (Cl−).
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, which is 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's ability to reduce is known as its reduction potential. The table below shows a few reduction potentials, which can be changed to oxidation potentials by reversing the sign. Reducing agents can be ranked by increasing strength by ranking their reduction potentials. Reducers donate electrons to (that is, "reduce") oxidizing agents, which are said to "be reduced by" the reducer. The reducing agent is stronger when it has a more negative reduction potential and weaker when it has a more positive reduction potential. The more positive the reduction potential the greater the species' affinity for electrons and tendency to be reduced (that is, to receive electrons). The following table provides the reduction potentials of the indicated reducing agent at 25 °C. For example, among sodium (Na), chromium (Cr), cuprous (Cu+) and chloride (Cl−), it is Na that is the strongest reducing agent while Cl− is the weakest; said differently, Na+ is the weakest oxidizing agent in this list while Cl is the strongest.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.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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