Covalent bonding

What Is a Covalent Bond?

A covalent bond is a shared pair of electrons between two non-metal atoms.

Each atom contributes one electron to the shared pair. The bond forms because the shared electrons are attracted to the nuclei of both atoms — this electrostatic attraction between the positive nuclei and the negative shared electrons holds the atoms together.

MisconceptionA covalent bond is not simply "sharing electrons." The examinable definition requires "a shared pair of electrons." One mark is often lost for omitting the word "pair."

Exam TipAlways write "shared pair of electrons" in definitions.

Dot-and-Cross Diagrams for Diatomic Molecules

In a dot-and-cross diagram, dots represent electrons from one atom and crosses represent electrons from the other. Only the outer-shell (valence) electrons are shown. A shared pair appears as one dot and one cross between the two atoms.

A double bond contains two shared pairs. A triple bond contains three shared pairs. Non-bonding electrons remain as lone pairs on the outer shell of the atom.

Dot-and-Cross Diagrams for Inorganic Molecules

• Water has two lone pairs on oxygen.
• Ammonia has one lone pair on nitrogen.
• Carbon dioxide has two C=O double bonds with no lone pairs on carbon.

Dot-and-Cross Diagrams for Organic Molecules

A structural formula writes atoms grouped by carbon: CH₃CH₃ means a chain of two carbons, each with hydrogens attached. In a displayed formula, every bond is drawn as a line between atoms.

Carbon always forms four bonds. Ethene contains a C=C double bond (two shared pairs between the two carbon atoms).

Simple Molecular Structures — Low Melting and Boiling Points

Substances with simple molecular structures are gases, liquids, or solids with low melting and boiling points.

Within each molecule, the covalent bonds are strong. However, between molecules there are weak intermolecular forces of attraction [forces acting between separate molecules, not within them]. When a simple molecular substance melts or boils, these weak intermolecular forces are overcome — not the covalent bonds. Because little energy is required to overcome weak intermolecular forces, the melting and boiling points are low.

MisconceptionStudents often state that "covalent bonds break" when a simple molecular substance boils. The covalent bonds within the molecules do not break — only the weak intermolecular forces between molecules are overcome.

Exam TipAlways specify "intermolecular forces are overcome" and state that covalent bonds are not broken.

Effect of Relative Molecular Mass on Boiling Point

The melting and boiling points of simple molecular substances increase, in general, with increasing relative molecular mass (Mᵣ).

Larger molecules have more electrons, which increases the strength of the intermolecular forces between them. Stronger intermolecular forces require more energy to overcome, so the boiling point increases.

Examiner InsightA common 3-mark question asks "Explain why bromine has a higher boiling point than chlorine." The required chain is: larger Mᵣ → stronger intermolecular forces → more energy needed to overcome them.

Exam TipAlways connect the three steps — mass, force strength, energy required.

Giant Covalent Structures — High Melting and Boiling Points

Substances with giant covalent structures are solids with very high melting and boiling points.

Every atom in a giant covalent structure is bonded to its neighbours by strong covalent bonds. These bonds extend throughout the entire structure in a continuous three-dimensional network. To melt or boil the substance, many strong covalent bonds must be broken. Energy is required to break bonds, so a very large amount of energy is needed — resulting in very high melting and boiling points.

MisconceptionIn giant covalent structures, covalent bonds are broken during melting. This is the opposite of simple molecular structures, where only intermolecular forces are overcome. Students sometimes apply the same reasoning to both.

Exam TipState clearly that "many strong covalent bonds must be broken" for giant covalent substances.

Diamond, Graphite, and C₆₀ Fullerene

Diamond, graphite, and C₆₀ fullerene are all forms of carbon with different structures that cause different physical properties.

Diamond — Each carbon atom forms four strong covalent bonds in a rigid three-dimensional tetrahedral network. No free electrons exist, so diamond does not conduct electricity. The strong bonding in all directions makes diamond extremely hard.

Graphite — Each carbon atom forms three covalent bonds, creating flat hexagonal layers. The fourth outer electron on each carbon is delocalised [free to move along the layers]. These delocalised electrons enable graphite to conduct electricity. The layers are held together by weak intermolecular forces, so layers slide over each other easily — making graphite soft and slippery.

C₆₀ fullerene — Sixty carbon atoms form a hollow spherical cage. Because C₆₀ molecules are individual units held together by weak intermolecular forces, fullerene has a much lower melting point than diamond or graphite. C₆₀ does not conduct electricity because there are no delocalised electrons free to move through the structure.

Examiner InsightQuestions on graphite require a precise three-part answer: (1) each carbon forms three covalent bonds, (2) one electron per carbon is delocalised, (3) delocalised electrons carry charge so graphite conducts electricity. Missing any part loses marks.

Exam TipPractise the three-part graphite conductivity answer until it is automatic.

Electrical Conductivity of Covalent Compounds

Covalent compounds do not usually conduct electricity.

In covalent substances, electrons are shared within bonds and are localised [held in fixed positions between atoms]. There are no free electrons or ions to carry charge. Therefore, simple molecular substances and most giant covalent substances cannot conduct electricity. Graphite is the notable exception because it contains delocalised electrons.