A reducing agent is a chemical species that donates electrons to an electron acceptor that is termed an oxidizing agent. In the process, the reducing agent is oxidized, while the oxidizing agent is reduced. Other names for a reducing agent are a reducer, reductant, or electron donor. Reducing agents and oxidizing agents always occur together in redox reactions. The word “redox” is a combination of the words “reduction” and “oxidation.” Examples of reducing agents include hydrogen gas, alkali metals, rare earth metals, and compounds containing the hydride (H–) anion.
- A reducing agent loses electrons and is oxidized in a chemical reaction. An oxidizing agent gains electrons an is reduced.
- Because both processes occur together, the reaction is a redox reaction.
- The oxidation state of a reducing agent increases in a redox reaction, while the oxidation state of an oxidizing agent decreases.
- Examples of reducing agents are hydrogen gas, group 1 and group 2 metals, and other reactants in low oxidation states.
Reducing Agent Word Origin
Originally, redox reactions involved the loss or gain of oxygen. An oxidizing agent gave its oxygen to the other species in the reaction, leaving it with a reduced amount of oxygen. The reducing agent reduced the amount of oxygen in the other species. Gaining oxygen made it oxidized.
How to Identify the Reducing Agent
But, redox reactions do not always involve oxygen. It’s all about the transfer of electrons, which changes the oxidation state.
- Reducing agents favor losing an electron to achieve a noble gas configuration.
- Oxidizing agent favor gaining an electron to achieve a noble gas configuration.
- Reducing agents are usually in a lower possible oxidation state.
- Oxidizing agents are usually in a higher possible oxidation state.
Identify a reducing agent (and oxidizing agent) by writing a balanced redox reaction and then separating it into balanced oxidation and reduction half-reactions.
For example, identify the reducing agent and oxidizing agent in this balanced equation for the reaction between chlorine and aqueous bromine ions:
Cl2(aq) + 2Br−(aq)⟶2Cl−(aq) + Br2 (aq)
In the balanced equation, bromine goes from the -1 oxidation state on the reactants side of the equation to the 0 oxidation state in the products side. Br– loses an electron. It is the reducing agent and is oxidized. Here is the oxidation half reaction:
2Br− (aq) ⟶ Br2 (aq)
Meanwhile, chlorine goes from the 0 oxidation state to the -1 oxidation state. It gains an electron, so it is the oxidizing agent and is reduced. Here is the reduction half reaction:
Cl2 (aq) ⟶ 2Cl− (aq)
Remembering Reducing Agents
Keeping reducing agent and oxidizing agent straight is confusing, but these chemistry mnemonics help:
- OIL RIG: Oxidation is loss of electrons; reduction is gain of electrons
- LEO (the lion) says GER: Loss of electrons is oxidation; gain of electrons is reduction
- LEORA says GEROA: This is similar to LEO says GER, except it includes reducing agent and oxidizing agent. The loss of electrons is oxidation (reducing agent), while the gain of electrons is reduction (oxidizing agent).
Examples of Reducing Agents
Here are examples of common commercial reducing agents. However, remember that the nature of the other species in the reaction matters! For example, sulfur dioxide acts as either a reducing agent or an oxidizing reagent, depending on the reaction.
- Hydrogen gas (H2)
- Iron(II) compounds (e.g., iron(II) sulfate)
- Tin(II) compounds (e.g., tin(II) chloride)
- Lithium aluminum hydride (LiAlH4)
- Red-Al [NaAlH2(OCH2CH2OCH3)2]
- Sodium amalgam (Na(Hg))
- Sodium-lead alloy (Na + Pb)
- Zinc amalgam [Zn(Hg)]
- Sodium borohydride (NaBH4)
- Sulfur dioxide (SO2, sometimes an oxidizing agent)
- Thiosulfates (e.g. Na2S2O3)
- Iodides (e.g., KI)
- Oxalic acid (C2H2O4)
- Formic acid (HCOOH)
- Ascorbic acid (C6H8O6)
- Carbon monoxide (CO)
- Carbon (C)
Can Oxygen Be a Reducing Agent?
Most of the time, oxygen is (as you might guess) an oxidizing agent. However, it can be a reducing agent. For example, in the reaction between oxygen and fluorine, oxygen is the reducing agent and fluorine is the oxidizing agent.
O2 (g) + 2 F2 (g) → 2 OF2 (g)
It’s easier seeing the process when you write the equation as half-reactions:
4 F + 4 e– → 4 F– (oxidizing agent, reduction)
2 O – 4 e– → 2 O2+ (reducing agent, oxidation)
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