Oxidation-Reduction Reactions

Oxidation-Reduction reactions involve transfer of electrons from one atom to another.

• The species which gives up electrons is oxidized.
• The species which accepts electrons is reduced.

To help remember this you may want to use the Oil rig mneumonic

Oxidation

Involves

Loss (of electrons)

Reduction

Involves

Gain (of electrons)

There are a wide variety of oxidation-reduction reactions, but as an introduction we are only going to consider oxidation of a metal.

M(s) + X ® Mⁿ⁺ + X^n-

The species X in the above equation takes electrons from the metal. It is called the oxidizing agent (because it causes the metal to be oxidized).

Consider reactions involving metals and the following three categories of oxidizing agents.

1. Oxygen

The classic oxidizing agent, molecular oxygen (O₂) removes electrons from the metal to form O^2- ions, and form a metal oxide product.

Examples

Al(s) + 3O₂ ® 2Al₂O₃(s)

Ti(s) + 2O₂ ® TiO₂(s)

2. Acids

We have learned that acids are a source of H⁺ ions, if you bring two H⁺ ions together and add two electrons you can form hydrogen gas (H₂). The source of electrons in this reaction is the metal.

Example

Fe(s) + H₂SO₄(aq) ® H₂(g) + FeSO₄(aq)

Net ionic equation

Fe(s) + 2H⁺(aq) ® H₂(g) + Fe²⁺(aq)

3. Other metal ions (in the form of soluble ionic compounds)

Some metals give up their electrons more easily than others. For example if you take a chunk of an alkali metal such as potassium and throw it into water you will get a vigorous, probably explosive reaction, as the potassium is oxidized to K⁺. On the other hand if you throw a piece of silver into water there will be no reaction. Therefore, we conclude that potassium gives up its electrons more readily than silver.

Now if we place K atoms in the presence of Ag⁺, the potassium is happy to give its electrons to the silver ions so that the silver ions can become neutral atoms again.

K(s) + AgNO₃(aq) ® KNO₃(aq) + Ag(s)

Net ionic equation

K(s) + Ag⁺(aq) ® K⁺(aq) + Ag(s)

The activity series is given in table 4.4 of your book. It ranks the various metals in order of their ease of oxidation. We can use the activity series to predict whether or not an oxidation-reduction reaction will occur. To do so simply remember the following rule:

Any metal in the activity series can be oxidized by the ions of the elements below it.

Examples

(a) 2Al(s) + 3NiCl₂(aq) ® 2AlCl₃(aq) + 3Ni(s)

Net ionic equation

2Al(s) + 3Ni²⁺(aq) ® 2Al³⁺(aq) + 3Ni(s)

(b) Ag(s) + Pb(NO₃)₂(aq) ® No reaction (N.R.)

Lead is not below silver in the activity series, so its ions cannot oxidize elemental silver and no reaction occurs.

(c) Mn(s) + 2HBr(aq) ® MnBr₂(aq) + H₂(g)

Net ionic equation

Mn(s) + 2H⁺(aq) ® Mn²⁺(aq) + H₂(g)

Note that the last example illustrates how we can use the activity series to predict which metals will be oxidized by an acid, and which ones will not react with acids (Cu, Ag, Hg, Pt, Au).