Structure & Nomenclature of Covalent Substances

Polar vs Non-polar Covalent Bonds

Polar and non-polar covalent bonds represent two extremes in the distribution of electron density in a covalent bond, and understanding the difference between them is crucial for predicting the behaviour and properties of molecules.

Polar Covalent Bonds

Polar covalent bonds occur between two atoms with differing electronegativities, leading to an unequal sharing of electrons. The more electronegative atom attracts the shared electron pair more strongly, acquiring a partial negative charge (δ⁻), while the less electronegative atom becomes partially positive (δ⁺). This separation of charge creates a dipole across the bond.

 

 

Examples of polar covalent bonds:

• Water (H₂O): The O-H bonds in water are polar because oxygen is significantly more electronegative than hydrogen. This causes a partial negative charge near the oxygen atom and a partial positive charge near the hydrogen atoms (a dipole), making the bond polar and water a polar molecule.

• Hydrogen Chloride (HCl): In an HCl molecule, chlorine is more electronegative than hydrogen, leading to a polar covalent bond. The chlorine atom bears a partial negative charge, and the hydrogen atom has a partial positive charge (dipole), resulting in a polar bond and molecule.

Non-polar Covalent Bonds

Non-polar covalent bonds occur when two atoms share a pair of electrons equally or close to equally. This usually happens between atoms of the same element or between different elements that have very similar electronegativities. Since the electron density is evenly distributed, there is no significant charge separation across the bond. Since there is no charge separation across a non-polar bond, there is no dipole present.

 

 

Examples of of non-polar covalent bonds:

• Diatomic Molecules: Molecules like H₂ (hydrogen gas), N₂ (nitrogen gas), O₂ (oxygen gas), and Cl₂ (chlorine gas) are classic examples of non-polar covalent bonds. Each molecule consists of two identical atoms sharing electrons equally. In diatomic molecules consisting of highly electronegative atoms e.g. O₂ and Cl₂, the covalent bonds are still non-polar because both atoms have equal attraction for electrons.

• Methane (CH₄): In methane (see diagram above), the carbon and hydrogen atoms have similar electronegativities, and the C-H bonds are considered non-polar. The symmetrical arrangement of the hydrogen atoms around the central carbon atom also ensures that any slight differences in electronegativity do not lead to a dipole.

Covalent Molecules

 

Covalent molecules consist of a discrete number of atoms bonded together covalently. These molecules can range from simple diatomic molecules like hydrogen (H₂) and simple covalent compounds like carbon dioxide (CO2) to large biomolecules like DNA.

 

Structure of deoxyribonucleic acid (DNA) molecules which are classified as covalent molecules.

Properties of Covalent Molecules:

• Low Melting and Boiling Points: Covalent molecules generally have lower melting and boiling points compared to ionic compounds, due to weaker intermolecular forces (like van der Waals forces) as opposed to the strong ionic bonds in ionic compounds.

• Poor Conductivity: In their pure state, covalent molecules do not conduct electricity, lacking free ions or electrons.

• Solubility: Nonpolar covalent molecules tend to be soluble in nonpolar solvents, while polar covalent molecules can dissolve in polar solvents, like water.

Physical properties of covalent molecules are related to intermolecular forces which is discussed separately here.

Covalent Network Substances

Covalent network substances, also known as network solids, are materials where atoms are bonded by covalent bonds in a continuous network extending throughout the material. Examples include diamond (a form of carbon) and silicon dioxide (SiO₂).

Figure (left): lattice structure of diamond consisting of carbon atoms connected by covalent bonds. Figure (right): lattice structure of silicon dioxide connected by covalent bonds.

Properties of Covalent Network Substances:

• High Melting and Boiling Points: The extensive network of strong covalent bonds gives these substances very high melting and boiling points.

• Hardness: Many covalent network substances are extremely hard, with diamond being the hardest known natural material.

• Poor Conductivity: Typically, network solids are poor conductors of electricity, as all electrons are localised in bonds. However, some network solids like graphite (another form of carbon) can conduct electricity due to the presence of delocalised electrons.

Similarities and Differences

While both covalent molecules and network substances are characterised by covalent bonds, their properties diverge significantly due to differences in structural organisation. The discrete nature of covalent molecules leads to weaker intermolecular forces and, consequently, lower melting and boiling points, whereas the continuous bonding in network substances results in much higher melting points and hardness.

Nomenclature of Covalent Substances

Naming covalent compounds involves using prefixes to indicate the number of atoms of each element present in the molecule. The prefixes are: mono- (1, often omitted for the first element), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8).

For example:

• CO₂ is carbon dioxide (not monocarbon dioxide)
• PCl₅ is phosphorus pentachloride
• N₂O₄ is dinitrogen tetroxide

Covalent vs. Ionic Naming

The main difference in naming covalent and ionic compounds lies in the use of prefixes for covalent compounds to denote the number of atoms, which is not done for ionic compounds. Ionic compounds are named by writing the cation name first followed by the anion name, with the anion typically ending in "-ide", such as in sodium chloride (NaCl). This difference reflects the distinct nature of ionic compounds, which consist of ions in a lattice structure and do not have discrete molecules, making the use of numerical prefixes unnecessary. 

The empirical formula of ionic compounds can be determined without using a prefix in their names by simply knowing the preferred oxidation states of metal and non-metal atoms.

For example, magnesium chloride: magnesium ion has an oxidation state of +2 (loses two electrons to reach complete valence shell), and chloride has an oxidation state –1 (gains one electron to reach complete valence shell). This means there must be two chlorine atoms for every one magnesium atom for them to both achieve a stable electron configuration.