Reducing agent

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 ferrocyanide ([Fe(CN)₆]⁴⁻). It donates an electron, becoming oxidized to ferricyanide ([Fe(CN)₆]³⁻). Simultaneously, the oxidizer chlorine is reduced to chloride.

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 reduce, or gain electrons, is known as its reduction potential.[1] 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. The reducing agent is stronger when it has a more negative reduction potential and weaker when it has a more positive reduction potential. The following table provides the reduction potentials of the indicated reducing agent at 25 °C. For example, among Na, Cr, Cu⁺ and Cl⁻, Na is the strongest reducing agent and Cl⁻ is the weakest one.

${\displaystyle {\begin{array}{|rl|r|}{\text{Oxidizing agent}}&\qquad {\text{Reducing agent}}&{\text{Reduction potential (V)}}\\\hline {\ce {{Li+}+e-}}&{\ce {<=>Li}}&-3.04\\{\ce {{Na+}+e-}}&{\ce {<=>Na}}&-2.71\\{\ce {{Mg^{2+}}+2e-}}&{\ce {<=>Mg}}&-2.38\\{\ce {{Al^{3+}}+3e-}}&{\ce {<=>Al}}&-1.66\\{\ce {{2H2O(l)}+2e-}}&{\ce {<=>{H2(g)}+2OH-}}&-0.83\\{\ce {{Cr^{3+}}+3e-}}&{\ce {<=>Cr}}&-0.74\\{\ce {{Fe^{2+}}+2e-}}&{\ce {<=>Fe}}&-0.44\\{\ce {{2H+}+2e-}}&{\ce {<=>H2}}&0.00\\{\ce {{Sn^{4+}}+2e-}}&{\ce {<=>Sn^{2+}}}&+0.15\\{\ce {{Cu^{2+}}+e-}}&{\ce {<=>Cu+}}&+0.16\\{\ce {{Ag+}+e-}}&{\ce {<=>Ag}}&+0.80\\{\ce {{Br2}+2e-}}&{\ce {<=>2Br-}}&+1.07\\{\ce {{Cl2}+2e-}}&{\ce {<=>2Cl-}}&+1.36\\{\ce {{MnO4^{-}}+{8H+}+5e-}}&{\ce {<=>{Mn^{2+}}+4H2O}}&+1.49\\{\ce {{F2}+2e-}}&{\ce {<=>2F-}}&+2.87\end{array}}}$ [2]

Common reducing agents include metals potassium, calcium, barium, sodium and magnesium, and also compounds that contain the H⁻ ion, those being NaH, LiH,[3] 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) + H_2(g) → 2 LiH_(s)[lower-alpha 1]

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

H_2(g) + F_2(g) → 2 HF_(g)[lower-alpha 2]

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

Reducing agents and oxidizing agents are the ones responsible for corrosion, which is the "degradation of metals as a result of electrochemical activity".[1] 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.

The formation of iron(III) oxide;

4Fe + 3O₂ → 4Fe³⁺ + 6O²⁻ → 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. 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).

• Lithium aluminium hydride (LiAlH₄), a very strong reducing agent
• Nascent (atomic) hydrogen
• 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)
• Compounds containing the Fe²⁺ ion, such as iron(II) sulfate
• Compounds containing 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, e.g. KI (mainly in analytical chemistry)
• Hydrogen peroxide (H
  ₂O
  ₂) – 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 (C
  ₂H
  ₂O
  ₄)
• Formic acid (HCOOH)
• Ascorbic acid (C₆H₈O₆)
• Reducing sugars
• Phosphites, hypophosphites, and phosphorous acid
• Dithiothreitol (DTT) – used in biochemistry labs to avoid S-S bonds
• Carbon monoxide (CO)
• Cyanides in hydrochemical metallurgical processes
• Carbon (C)
• Tris-2-carboxyethylphosphine hydrochloride (TCEP)

• Organic reduction
• Oxidizing agent
• Electrochemistry
• Electrosynthesis
• Corrosion
• Electrolyte
• Redox
• Reducing equivalent

• "Chemical Principles: The Quest for Insight", Third Edition. Peter Atkins and Loretta Jones p. F76

• Table summarizing strength of reducing agents at the Wayback Machine (archived June 11, 2011)