Structure & Nomenclature of Covalent Substances


This is part of Year 11 HSC Chemistry course under the topic of Bonding.

HSC Chemistry Syllabus

    • Investigate the different chemical structures of atoms and elements, including but not limited to:
    – Covalent networks (including diamond and silicon dioxide)
    – Covalent molecular
    • investigate the role of electronegativity in determining the ionic or covalent nature of bonds between atoms

        Structure, Properties and Nomenclature of Covalent Substances

        What is a Covalent Bond?


        Covalent bonds form when two atoms share one or more pairs of valence electrons, stabilising each other by filling their outer electron shells (octet rule). This type of bonding is common between non-metal atoms, which have high electronegativities and a strong tendency to attract electrons. The shared electron pairs are referred to as bonding pairs, while the unshared pairs associated with each atom are called lone pairs.

        For example, a covalent bond can be formed between two atoms of hydrogen to achieve a complete shell (2 electrons) for both atoms.

        In a molecule of water, the oxygen atom forms two covalent bonds with the two hydrogen atoms to achieve a complete valence shell (8 electrons). These two covalent bonds also help the two hydrogen atoms achieve a stable configuration.


        Covalent bonds manifest in three distinct forms: single, double, and triple bonds.

        • Single Bond: A single bond forms when a pair of electrons is shared between two atoms. It is symbolised by a single line linking the two atoms. 

        • Double Bond: A double bond arises when two electron pairs are shared between two atoms, represented by two parallel lines between the atoms in a molecule. An example of a molecule containing a double covalent bond is carbon dioxide (valency of oxygen is 2, and carbon is 4).

        • Triple Bond: A triple bond is formed when three pairs of electrons are shared between two atoms in a molecule. An example of a molecule containing a triple covalent bond is nitrogen gas (valency of nitrogen is 3).

        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.


        Polar bond in water


        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.


        Non-polar covalent bond in methane molecule


        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


        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 DNA bases

        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.