A simple redox reaction occurs when copper metal is immersed in a solution of silver nitrate. The solution gradually acquires the blue color characteristic of the hydrated Cu²⁺ ion, while the copper becomes coated with glittering silver crystals. The reaction may be described by the net ionic equation

Cu(s) + 2Ag⁺(aq) ? Cu²⁺(aq) + 2Ag(s)      (1)

We can regard this equation as being made up from two hypothetical half-equations. In one, each copper atom loses 2 electrons:

Cu ? Cu²⁺ + 2e^–      (1a)

while in the other, 2 electrons are acquired by 2 silver ions:

2e^– + 2Ag⁺ ? 2Ag      (1b)

If these two half-equations are added, the net result is Eq. (1). In other words, the reaction of copper with silver ions, described by Eq. (1), corresponds to the loss of electrons by the copper metal, as described by half- equation (1a), and the gain of electrons by silver ions, as described by Eq. (1b).

A species like copper which donates electrons in a redox reaction is called a reducing agent, or reductant. (A mnemonic for remembering this is remember, electron donor = reducing agent.) When a reducing agent donates electrons to another species, it is said to reduce the species to which the electrons are donated. In Eq. (1), for example, copper reduces the silver ion to silver. Consequently the half-equation

2Ag⁺ + 2e^– ? 2Ag

is said to describe the reduction of silver ions to silver.

Species which accept electrons in a redox reaction are called oxidizing agents, or oxidants. In Eq. (1) the silver ion, Ag⁺, is the oxidizing agent. When an oxidizing agent accepts electrons from another species, it is said to oxidize that species, and the process of electron removal is called oxidation. The half-equation

Cu ? Cu²⁺ + 2e^–

thus describes the oxidation of copper to Cu²⁺ ion.

In summary, then, when a redox reaction occurs and electrons are transferred, there is always a reducing agent donating electrons and an oxidizing agent to receive them. The reducing agent, because it loses electrons, is said to be oxidized. The oxidizing agent, because it gains electrons, is said to be reduced.

Copper is also oxidized by the oxygen present in air. The following video shows an example of this oxidation occurring. It also demonstrates how this oxidation can be reversed by using hydrogen gas.

EXAMPLE 1 Write the following reaction in the form of half-equations. Identify each half-equation as an oxidation or a reduction. Also identify the oxidizing agent and the reducing agent in the overall reaction

Zn + 2Fe³⁺ ? Zn²⁺ + 2Fe²⁺

Worked Problem Here

A more complex redox reaction occurs when copper dissolves in nitric acid. The acid attacks the metal vigorously, and large quantities of the red-brown gas, nitrogen dioxide (NO₂) are evolved. (NO₂ is poisonous, and so this reaction should be done in a hood.) The solution acquires the blue color characteristic of the hydrated Cu²⁺ ion. The reaction which occurs is
Cu + 2NO₃^– + 4H₃O⁺ ? Cu²⁺ + 2NO₂ + 6H₂O      (2)

Merely by inspecting this net ionic equation, it is difficult to see that a transfer of electrons has occurred. Clearly the copper metal has lost electrons and been oxidized to Cu²⁺, but where have the donated electrons gone? The matter becomes somewhat clearer if we break up Eq. (2) into half- equations. One must be

Cu ? Cu²⁺ + 2e^–      (2a)

and the other is

2e^– + 4H₃O⁺ + 2NO₃^– + ? 2NO₂ + 6H₂O      (2b)

You can verify that these are correct by summing them to obtain Eq. (2).

The second half-equation shows that each NO₃^– ion has not only accepted an electron, but it has also accepted two protons. To further complicate matters, a nitrogen-oxygen bond has also been broken, producing a water molecule. With all this reshuffling of nuclei and electrons, it is difficult to say whether the two electrons donated by the copper ended up on an NO₂ molecule or on an H₂O molecule. Nevertheless, it is still meaningful to call this a redox reaction. Clearly, copper atoms have lost electrons, while a combination of hydronium ions and nitrate ions have accepted them. Accordingly, we can refer to the nitrate ion (or nitric acid, HNO₃) as the oxidizing agent in the overall reaction. Half-equation (2b) is a reduction because electrons are accepted. 

To determine whether a Redox reaction has taken place you often must write the oxidation states of the species involved. You learned this method previously in the first semester of general chemistry, but the following page will provide you with a review. 

Subpages (2): Example 1 Oxidation Numbers Review