A reducing agent is an element or compound that loses an electron to another chemical
species in a redox chemical reaction. Since the reducing agent is losing electrons, it
is said to have been oxidized. If any chemical is an electron donor, another
must be an electron recipient. A reducing agent is oxidized because it loses electrons
in the redox reaction. Thus reducers are "oxidized" by oxidizers and oxidizers are "reduced" by
reducers; reducers are by themselves reduced and oxidizers are by themselves oxidized.
A reducing agent typically is in one of its lower possible oxidation states and is known
as the electron donor. Examples of reducing agents include the earth metals, formic acid,
and sulfite compounds. For example, consider the overall reaction
for aerobic cellular respiration: C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l)
The oxygen is being reduced, so it is the oxidizing agent. The glucose is being oxidized,
so it is the reducing agent. 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.
Historically, reduction referred to the removal of oxygen from a compound, hence the name
'reduction'. The modern sense of donating electrons is a generalisation of this idea,
acknowledging that other components can play a similar chemical role to oxygen.
Characteristics Consider the following reaction:
2 [Fe(CN)6]4− + Cl 2 → 2 [Fe(CN)6]3− + 2 Cl−
The reducing agent in this reaction is ferrocyanide6]4-). It donates an electron, becoming oxidized
to ferricyanide6]3-). Simultaneously, the oxidizer chlorine is reduced to chloride.
Strong reducing agents easily lose 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
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, thus oxidation always
occurs in the anode, and the cathode is an element that gains electrons, 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 → 2Fe23+O36- In the above equation, the Iron has an oxidation
number of 0 before and 3+ after the reaction. For oxygen 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 has been oxidized because the oxidation number increased. Iron is the reducing agent
because it gave electrons to the oxygen. Oxygen has been reduced because the oxidation number
has decreased and is the oxidizing agent because it took electrons from iron. ferric
Common reducing agents Lithium aluminum hydride
Nascent hydrogen Sodium amalgam
Diborane Sodium borohydride
Compounds containing the Sn2+ ion, such as tin(II) chloride
Sulfite compounds Hydrazine
Zinc-mercury amalgam) Diisobutylaluminum hydride
Lindlar catalyst Oxalic acid
Formic acid Ascorbic acid
Phosphites, hypophosphites, and phosphorous acid
Dithiothreitol – used in biochemistry labs to avoid S-S bonds
Compounds containing the Fe2+ ion, such as iron(II) sulfate
Carbon monoxide(CO) Carbon
Tris(2-carboxyethyl)phosphine HCl Several common reducing agents and their products
See also Organic reduction
Oxidizing agent Electrochemistry
Electrosynthesis Corrosion
Electrolyte Redox
References
Further reading "Chemical Principles: The Quest for Insight",
Third Edition. Peter Atkins and Loretta Jones p. F76
External links Table summarizing strength of reducing agents
at the Wayback Machine