## Common Oxidizing and Reducing Agents

Read these sections, which look at some common oxidizing and reducing agents, and substances that can act as an oxidizing agent or reducing agent depending on what they react with. Water is one of the compounds that can act as an oxidizing or reducing agent.

### 1. Common Oxidizing Agents

Oxidizing agents must be able to accept electrons readily. Elements with high electronegativity readily accept electrons, as can molecules or ions which contain relatively electronegative elements and even some metals which have high oxidation numbers. Bear these general rules in mind as we examine examples of common oxidizing agents in the following paragraphs.

### Oxidizing Agents

#### Halogens (group VllA elements)

All four elemental halogens, F2, Cl2, Br2, and I2, are able to accept electrons according to the half-equation

$X_{2}+2e^{-}\rightarrow 2X^{-}$

with X = F, Cl, Br, I

As we might expect from the periodic variation of electronegativity, the oxidizing power of the halogens decreases in the order F2 > Cl2 > Br2 > I2. Fluorine is such a strong oxidizing agent that it can react with water (water is very difficult to oxidize):

$2F_{2}+6H_{2}O\rightarrow 4H_{3}O^{+}+4F^{-}+O_{2}$

Chlorine also reacts with water, but only in the presence of sunlight. Bromine is weaker, and iodine has only mild oxidizing power.

#### Oxygen

Oxygen gas, which constitutes about 20 percent of the earth’s atmosphere, is another electronegative element which is a good oxidizing agent. It is slightly weaker than chlorine, but considerably stronger than bromine. Because the atmosphere contains such a strong oxidant, few substances occur in reduced form at the earth’s surface.

An oxidized form of silicon, SiO2, is one of the most plentiful constituents of the crust of the earth. Most metals, too, occur as oxides and must be reduced before they can be obtained in elemental form. When iron rusts, it forms the red-brown oxide Fe2O3n x H2O, seen below, which always contains an indeterminate amount of water.

Iron Oxide (Fe2O3 x H2O)

#### Oxyanions and Oxyacids

In aqueous solution NO3, IO3, MnO4, Cr2O72–, and a number of other oxyanions serve as convenient, strong oxidizing agents. The structure of the last oxyanion mentioned above is shown in Figure 1. The most strongly oxidizing oxyanions often contain an element in its highest possible oxidation state, that is, with an oxidation number equal to the periodic group number. For example, NO3 contains nitrogen in a +5 oxidation state, Cr2O72– (seen below) contains chromium +6, and has manganese +7.

The oxidizing power of the dichromate ion is employed in laboratory cleaning solution, a solution of Na2Cr2O7 in concentrated H2SO4. This readily oxidizes the organic compounds in grease to carbon dioxide. It is also highly corrosive, eats holes in clothing, and must be handled with care. Dark purple permanganate ion is another very common oxidizing agent (seen below). In basic solution it is reduced to solid dark brown MnO2. In acidic solution, however, it forms almost colorless Mn2+(aq).

Space-filling (left) and ball-and-stick models (right) of dichromate ion, Cr2O72–. Chromium atoms are gray and oxygen atoms are dark red.

A beaker of potassium permanganate solution, which is a very dark purple even at low concentrations. Image: David Mülheims CC BY SA 2.5, via Wikimedia Commons.

### 2. Common Reducing Agents

A good reducing agent must be able to donate electrons readily, meaning it must not have a high electronegativity. Among the elements, low electronegativity is characteristic of good reducing agents. Molecules and ions which contain relatively electropositive elements which have low oxidation numbers are also good reducing agents. Bear these general rules in mind as we examine examples of common reducing agents in the following paragraphs.

### Reducing Agents

#### Metals

All metals have low ionization energies and are relatively electropositive, and so they lose electrons fairly easily. Therefore, most metals are good reducing agents. Metals on the left of the periodic table exhibit this property to the greatest extent, and some of them, such as Li or Na, can even reduce H2O:

$2Li(s)+2H_{2}O(l)\rightarrow Li^{+}(aq)+2OH^{-}(aq)+H_{2}(g)$

Other metals, such as Fe or Zn, cannot reduce H 2O but can reduce hydronium ions, and so they dissolve in acid solution:

$Zn(s)+2H_{3}O^{+}(aq)\rightarrow Zn^{2+}(aq)+2H_{2}O(l)+H_{2}(g)$

This is one of the characteristic reactions of acids. There are a few metals that will not dissolve in just any acid but instead require an acid like HNO3 whose anion is a good oxidizing agent. Cu and Hg are examples:

$3Hg(s)+8H_{3}O^{+}(aq)+2NO_{3}^{-}(aq)\rightarrow 3Hg^{2}(aq)+NO(g)+12H_{2}O$

Finally, a few metals, such as Au and Pt, are such poor reducing agents that even an oxidizing acid like HNO3 will not dissolve them. This is the origin of the phrase the acid test. If a sample of an unknown yellow metal can be dissolved in acid, then the metal is not gold. Kings who collected tax payments in gold kept a supply of HNO3 available to make sure they were not being cheated.

### 3. Substances which are Both Oxidizing and Reducing Agents

In the section on acids and bases, we saw that some substances can act as both an acid and a base (amphiprotic). In the world of redox chemistry there exist substances that can act as both a reducing agent and oxidizing and a couple of examples are given below.

#### Water

We have seen that some oxidizing agents, such as fluorine, can oxidize water to oxygen. There are also some reducing agents, such as lithium, which can reduce water to hydrogen. In terms of redox, water behaves much as it did in acid-base reactions, where we found it to be amphiprotic.

In the presence of a strong electron donor (strong reducing agent), water serves as an oxidizing agent. In the presence of a strong electron acceptor (strong oxidizing agent), water serves as a reducing agent. Water is rather weak as an oxidizing or as a reducing agent, however; so there are not many substances which reduce or oxidize it. Thus it makes a good solvent for redox reactions. This also parallels water’s acid-base behavior, since it is also a very weak acid and a very weak base.

#### Hydrogen Peroxide (H2O2)

In this molecule the oxidation number for oxygen is –1. This is halfway between O2(0) and H2O(–2), and so hydrogen peroxide can either be reduced or oxidized.

When it is reduced, it acts as an oxidizing agent:

$H_{2}O_{2}+2H^{+}+2e^{-}\rightarrow 2H_{2}O$

When it is oxidized, it serves as a reducing agent:

$H_{2}O_{2}\rightarrow O_{2}+2H^{+}+2e^{-}$

Hydrogen peroxide is considerably stronger as an oxidizing agent than as a reducing agent, especially in acidic solutions.