10.4 - Metal Oxides & Non-metal Oxides - notes

Lewis acids and bases

Lewis acids and bases are defined by their ability to accept or donate electron pairs:

• A Lewis acid can accept a pair of electrons.
• A Lewis base can donate a pair of electrons.

Many metal oxides behave as Lewis bases, while non-metal oxides act as Lewis acids.

Metal oxides as Lewis bases

Metal oxides, particularly those of alkali and alkaline earth metals, act as Lewis bases. They donate an electron pair to hydrogen when reacting with water, forming hydroxides.

Alkali metal oxides react with water according to the general equation:

M₂O_(s) + H₂O_(l) ➔ 2MOH_(aq)

Examples include:

Li₂O_(s) + H₂O_(l) ➔ 2LiOH_(aq)

Na₂O_(s) + H₂O_(l) ➔ 2NaOH_(aq)

Group 2 metal oxides also react with water to produce hydroxides:

MO_(s) + H₂O_(l) ➔ M(OH)_2(aq)

Examples include:

MgO_(s) + H₂O_(l) ➔ Mg(OH)_2(aq)

CaO_(s) + H₂O_(l) ➔ Ca(OH)_2(aq)

Non-metal oxides as Lewis acids

Non-metal oxides behave as Lewis acids, accepting an electron pair from oxygen when reacting with water to form acidic compounds.

Some examples of non-metal oxides reacting with water include:

CO_2(g) + H₂O_(l) ➔ H₂CO_3(aq)

SO_2(g) + H₂O_(l) ➔ H₂SO_3(aq)

SO_3(l) + H₂O_(l) ➔ H₂SO_4(aq)

P₄O_10(s) + 6H₂O_(l) ➔ 4H₃PO_4(aq)

Amphoteric oxides

Some oxides, such as aluminium oxide (Al₂O₃), display both Lewis acidic and basic properties depending on the reaction partner. These oxides are called amphoteric oxides.

When reacting with a base like sodium hydroxide, Al₂O₃ behaves as a Lewis acid:

Al₂O_3(s) + 2NaOH_(aq) + 3H₂O_(l) ➔ 2NaAl(OH)_4(aq)

On the other hand, when reacting with an acid such as hydrochloric acid, Al₂O₃ acts as a Lewis base:

Al₂O_3(s) + 6HCl_(aq) ➔ 2AlCl_3(aq) + 3H₂O_(l)

Both of these reactions are examples of neutralisation, in which an acid and a base react to form a salt.

Trends in acid-base properties across a period

As you move from left to right across a period in the periodic table, the oxides of the elements exhibit a gradual change from basic to amphoteric to acidic. This trend is a consequence of the increasing covalent character of the bonds in the oxides.

• Metal oxides on the left side of the period (e.g., Na₂O and MgO) have primarily ionic bonding, which results in their basic properties when reacting with water.
• Oxides of elements in the middle of the period, like Al₂O₃, display both ionic and covalent bonding, leading to their amphoteric behaviour.
• Non-metal oxides on the right side of the period (e.g., P₄O₁₀, SO₃, and SO₂) have covalent bonding, which accounts for their acidic properties in water.

This gradual change in acid-base properties across a period can be summarised as follows:

Basic metal oxides ➔ Amphoteric oxides ➔ Acidic non-metal oxides

The table below illustrates the trend in acid-base properties of some period 3 oxides:

Acid rain

Acid rain is rainwater with a pH below 5.6, caused by the dissolution of nitrogen oxides (NO_x) and sulfur dioxide (SO₂) in the atmosphere. These gases can be released naturally from volcanoes and decaying vegetation or from human activities such as burning high-sulfur fossil fuels in industrial processes.

When sulfur dioxide dissolves in rainwater, it forms sulfurous acid:

SO_2(g) + H₂O_(l) ➔ H₂SO_3(aq)

Similarly, when nitrogen dioxide dissolves in rainwater, it forms nitric and nitrous acids:

2NO_2(g) + H₂O_(l) ➔ HNO_3(aq) + HNO_2(aq)

Atmospheric oxygen can further oxidise sulfurous and nitrous acids to sulfuric and nitric acids, respectively:

2H₂SO_3(aq) + O_2(g) ➔ 2H₂SO_4(aq)

2HNO_2(aq) + O_2(g) ➔ 2HNO_3(aq)

The increased acidity of acid rain can harm plants, aquatic life, and infrastructure.

Ocean acidification

Ocean acidification is the decrease in the pH of the Earth's oceans due to the absorption of large amounts of atmospheric carbon dioxide (CO₂).

When CO₂ dissolves in seawater, it forms carbonic acid:

CO_2(g) + H₂O_(l) ➔ H₂CO_3(aq)

The formation of carbonic acid lowers the pH of the oceans, making them less alkaline and more acidic. As human activities continue to increase atmospheric CO_2_ levels, more CO_2_ is absorbed by the oceans, leading to a further decrease in pH.

Ocean acidification can have severe consequences for marine life, especially organisms that build calcium carbonate skeletons or shells, such as coral reefs and shellfish. The lower pH makes it more challenging for these organisms to construct and maintain their structures, as the acidic conditions can cause the dissolution of calcium carbonate.