A reducing agent (also called a reductant or reducer) is the element or a compound in a redox (reduction-oxidation) reaction (see electrochemistry) that reduces another species. In doing so, it becomes oxidized, and is therefore the electron donor in the redox. For example consider the following reaction:
- [Fe(CN)6]4- + 1/2 Cl2 → [Fe(CN)6]3- + Cl
- C6H6 + 3 H2 → C6H12
In organic chemistry, good reducing agents are reagents that deliver H2.
What makes a strong reducing agent?
Strong reducing agents easily lose (or donate) electrons. Atoms with relatively large atomic radii tend to be better reductants. 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 electrons1, and 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"2. The table below shows a few reduction potentials that could easily be changed to oxidation potential by simply changing the sign. Reducing agents can be ranked by increasing strength by ranking their oxidation potentials. The reducing agent will be the strongest when it has a more positive oxidation potential and will be a weak reducing agent whenever it has a negative oxidation potential. The following table provides the reduction potentials of the indicated reducing agent at 25° C. Also remember the useful mnemonic devices, "OIL RIG," which means Oxidation Is Loss (of electrons) and Reduction Is Gain (of electrons), or "LEO the lion says GER," which means Loss of Electrons is Oxidation and Gain of Electrons is Reduction.
|Oxidizing Agent||Reducing Agent||Reduction Potential (v)|
|Li+ + e- =||Li||-3.04|
|Na+ + e- =||Na||-2.71|
|Mg2+ + 2e- =||Mg||-2.38|
|Al3+ + 3e- =||Al||-1.66|
|2H2O(l) + 2e- =||H2(g) + 2OH -||-0.83|
|Cr3+ + 3e- =||Cr||-0.74|
|Fe2+ + 2e- =||Fe||-0.41|
|2H+ + e- =||H2||0.00|
|Sn4+ + 2e- =||Sn2+||+0.15|
|Cu2+ + e- =||Cu+||+0.16|
|Ag+ + e- =||Ag||+0.80|
|Br2 + 2e- =||2Br-||+1.07|
|Cl2 + 2e- =||2Cl-||+1.36|
|MnO4-2 + 8H+ + 5e- =||Mn+2 + 4H2O||+1.49|
In order to tell which is the strongest reducing agent, change the sign of its respective reduction potential in order to make it oxidation potential. The bigger the number the stronger a reducing agent it is.
For example if one were to list Cu, Cl-, Na and Cr in order, one would get their reduction potential, change the sign to make it oxidation potential and list them from greatest to least. One will get Na, Cr, Cu and Cl-; Na being the strongest reducing agent and Cl- being the weakest one.
A few good common reducing agents include active metals such as potassium, calcium, barium, sodium and magnesium and also, compounds that contain the H- ion, those being NaH, LiAlH4 and CaH2.
Also, 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.
2Li(s) + H2(g) -->2LiH(s) hydrogen acts as an oxidizing agent because it accepts an electron donation from lithium, which causes Li to be oxidized.
Half Reactions 2Li(s)0 -->2Li(s)+1 + 2e-::::: H20(g) + 2e- --> 2H-1(g)
H2(g) + F2(g) --> 2HF(g) hydrogen acts as a reducing agent because it donates its electrons to fluorine, which allows fluorine to be reduced.
Half Reactions H20(g) --> 2H+1(g) + 2e-::::: F20(g) + 2e- --> 2F-1(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”3. 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 will begin deteriorating given that 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:
- Oxidation Half Reaction: Fe0 → Fe3+ + 3e−
- Reduction Half Reaction: O2 + 4e− → 2 O2−
Iron III (Fe) has been oxidized because the oxidation number increased and is the reducing agent because it gave electrons to the oxygen (O). Oxygen (O) has been reduced because the oxidation number has decreased and is the oxidizing agent because it took electrons from iron (Fe)
Common reducing agents
- Ferrous ion
- Lithium aluminium hydride (LiAlH4)
- Nascent hydrogen
- Sodium amalgam
- Sodium borohydride (NaBH4)
- Stannous ion
- Sulfite compounds
- Hydrazine (Wolff-Kishner reduction)
- Zinc-mercury amalgam (Zn(Hg)) (Clemmensen reduction)
- Diisobutylaluminum hydride (DIBAH)
- Lindlar catalyst
- Oxalic acid (C2H2O4)
Common reducing agents and their products
|H2 Hydrogen||H+, H2O|
|C||CO2 carbon dioxide|
|hydrocarbons||CO2 carbon dioxide, H2O|
- "Chemical Principles: The Quest for Insight", Third Edition. Peter Atkins and Loretta Jones pg. F76