Chemical Bonding and Molecular Geometry

Bonding and Structure: The Basics

You can snap a piece of pencil graphite with your fingers, but try snapping a diamond — even though both are made entirely of carbon! The secret to a material's physical behaviour lies completely in how its atoms bond and arrange themselves.

Ionic Bonding: Transferring Electrons

An ionic bond forms between a metal and a non-metal when atoms transfer electrons to achieve a full outer shell (matching the stable electronic structure of a noble gas). First, metal atoms lose electrons from their outer shell to become positively charged ions (cations). Then, non-metal atoms gain these transferred electrons to become negatively charged ions (anions). For example, in sodium chloride ($NaCl$), sodium (Group 1) undergoes electron transfer by giving its 1 outer electron to chlorine (Group 7). This creates a $Na^+$ ion and a $Cl^-$ ion. These oppositely charged ions are held together by extremely strong electrostatic forces of attraction. These forces act in all directions to form a giant ionic lattice — a regular, 3D arrangement of alternating positive and negative ions.

Covalent Bonding: Sharing Electrons

When two non-metals react, they do not transfer electrons; instead, they form a covalent bond by sharing pairs of electrons in their overlapping outer shells. The covalent bond is the strong electrostatic attraction between the shared pair of electrons and the positive nuclei of both atoms. By sharing electrons, both atoms achieve a stable, full outer shell. Atoms can share one pair to form a single bond (like in methane, $CH_4$, or water, $H_2O$). They can also share multiple pairs.

For example, the formation of an oxygen molecule ($O_2$) happens step-by-step:

1. Oxygen is in Group 6, meaning each atom has 6 electrons in its outer shell.
2. Each oxygen atom needs 2 more electrons to achieve a full outer shell (the stable structure of a noble gas).
3. Two oxygen atoms overlap their outer shells and share two pairs of electrons.
4. This results in a double bond ($O=O$), where 4 electrons in total are shared in the overlap, giving both atoms a stable, full outer shell.

Similarly, nitrogen ($N_2$) is in Group 5 and needs 3 electrons, so two nitrogen atoms overlap their outer shells and share three pairs of electrons. This forms a triple bond ($N\equiv N$).

Metallic Bonding and Properties

Metals consist of a giant lattice of positive metal ions arranged in a regular pattern. Because the outer electrons of metal atoms are shed, they form a "sea" of delocalised electrons that are free to move throughout the structure. The strong electrostatic attraction between the positive ions and these negative electrons holds the metal together. Metals are malleable (can be hammered into shape) and ductile (can be drawn into wires) because the metal ions are arranged in regular layers. When a force is applied, these layers can simply slide over each other while the mobile electrons hold the structure together.

Explaining Melting and Boiling Points

To understand why substances melt or boil at different temperatures, you must look at their structure and the specific forces you are trying to overcome. In a giant covalent structure (like diamond or silicon dioxide, $SiO_2$), billions of atoms are linked by strong covalent bonds in a continuous 3D lattice. Melting these requires breaking many strong covalent bonds, requiring massive amounts of energy and resulting in very high melting points. Similarly, ionic compounds have high melting points because large amounts of energy are needed to overcome the strong electrostatic forces acting in all directions across the giant ionic lattice. However, a simple molecular structure (like $Cl_2$, $H_2O$, or fullerenes like $C_{60}$) behaves differently. They consist of a fixed, finite number of atoms with strong covalent bonds inside the molecule, but only weak intermolecular forces between separate molecules. When these simple molecules boil or melt, you only overcome the weak intermolecular forces — the strong covalent bonds remain intact. This requires low energy, resulting in low melting and boiling points. As molecules get larger (like moving down Group 7 or in long chain polymers), the intermolecular forces become stronger, requiring more energy to overcome and raising the boiling point.

Comparing Giant and Simple Structures

When explaining physical properties in exams, comparing the forces at play is essential:

Feature	Giant Covalent Structure	Simple Molecular Structure
Structure Size	Billions of atoms in a continuous 3D lattice	Fixed, finite number of atoms in a molecule
Forces to Overcome	Strong covalent bonds must be broken	Weak intermolecular forces must be overcome
Energy Required	Massive amounts of energy	Very low amounts of energy
Melting/Boiling Point	Very high (e.g., diamond >3500°C)	Low (usually gases or liquids at room temperature)

Explaining Electrical Conductivity

For a substance to conduct electricity, it must contain charged particles (either ions or electrons) that are free to move and carry the charge. Solid ionic compounds act as insulators because their ions are locked in fixed positions within the lattice. However, when melted (molten) or dissolved in water (aqueous), the lattice breaks apart and the ions are free to move, allowing them to conduct electricity. Metals are excellent conductors in both solid and liquid states because their delocalised electrons are constantly free to move through the structure. Simple molecules and most giant covalent structures (like diamond) are insulators because they have no mobile charge carriers. The major exception is graphite: each carbon atom only forms 3 bonds, leaving a delocalised electron that can carry charge between its flat, hexagonal layers.