How Chemical Bonds Hold Atoms Together in Molecules

A Bond Stronger Than Steel Holds Carbon and Hydrogen Together

The carbon-hydrogen bond in methane has a bond dissociation energy of 439 kJ/mol — meaning it takes 439,000 joules to break one mole of these bonds. Yet this bond forms because the bonded state is lower in energy than the separated atoms. Chemical bonds are not physical tethers; they are regions of lowered energy arising from the electrostatic interactions between nuclei and electrons. Understanding them requires quantum mechanics, but their practical consequences shape every material property from diamond's hardness to protein folding in living cells.

Three primary types of chemical bonds — covalent, ionic, and metallic — along with weaker intermolecular interactions, determine how atoms assemble into the molecules and materials of the observable world.

Covalent Bonds

Covalent bonds form when atoms share pairs of electrons, lowering the total energy of the system compared to separated atoms. The bond arises because shared electrons interact attractively with both nuclei simultaneously, overcoming the nucleus-nucleus repulsion.

• A single bond shares one electron pair (bond order = 1); a double bond shares two pairs (bond order = 2); a triple bond shares three (bond order = 3).
• Higher bond order means shorter, stronger bonds: carbon-carbon single bond is 154 pm and 346 kJ/mol; triple bond is 120 pm and 835 kJ/mol.
• Molecular orbital theory describes bonding more precisely: atomic orbitals combine to form bonding and antibonding molecular orbitals. Electrons in bonding orbitals lower energy; electrons in antibonding orbitals raise it.
• Bond polarity arises from electronegativity differences: if one atom pulls electrons more strongly, the bond has a partial negative charge on one end and partial positive on the other, denoted δ− and δ+.

Ionic Bonds

Ionic bonds form through electron transfer between atoms of very different electronegativity. A metal with low ionisation energy donates electrons to a non-metal with high electron affinity. The resulting cations and anions attract electrostatically in a three-dimensional lattice.

Sodium chloride is the textbook example. Sodium's first ionisation energy is 496 kJ/mol; chlorine's electron affinity is 349 kJ/mol. The lattice energy of NaCl — the energy released when gaseous Na⁺ and Cl⁻ ions form the crystal — is approximately 788 kJ/mol, making the overall process highly exothermic. Lattice energy depends on ion charge and size: magnesium oxide (Mg²⁺ and O²⁻) has a lattice energy of about 3,850 kJ/mol and a melting point of 2,852°C.

• Ionic compounds are hard but brittle: the lattice resists compression but shatters when planes of ions shift and like charges face each other.
• Ionic compounds dissolve in polar solvents like water: the polar water molecules surround and stabilise individual ions (solvation), overcoming the lattice energy.
• Molten ionic compounds conduct electricity because ions are mobile; solid ionic compounds do not.

Metallic Bonds

Metallic bonding arises in metals where outer electrons are delocalised — not bound to individual atoms but free to move through the entire crystal lattice. Metal atoms exist as cations in a sea of mobile electrons.

Hydrogen Bonds and Intermolecular Forces

Weaker than covalent bonds but critical to biology and materials science, intermolecular forces arise between molecules rather than within them.

The hydrogen bond forms between a hydrogen atom covalently bonded to a highly electronegative atom (N, O, or F) and a lone pair on another electronegative atom. The N–H···O hydrogen bond has an energy of about 20 kJ/mol — roughly 20 times weaker than a covalent bond but far stronger than other intermolecular forces.

• Water's anomalously high boiling point (100°C vs −60°C expected from molecular mass trends) arises from its four hydrogen bonds per molecule in liquid form.
• DNA's double helix is held together by hydrogen bonds between complementary bases — adenine pairs with thymine via two hydrogen bonds, guanine with cytosine via three.
• Protein secondary structure (alpha helices, beta sheets) depends entirely on backbone hydrogen bonds.
• Van der Waals forces — induced dipole-dipole interactions — scale with molecular size and allow geckos to cling to glass via billions of tiny hairs engaging London dispersion forces.

Molecular Shape and VSEPR

Bond type determines which atoms combine; bond angles and molecular geometry determine the shape of the resulting molecule. The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts geometry from the number of electron pairs around the central atom: electron pairs repel one another and adopt arrangements that maximise their separation.

Water has two bonding pairs and two lone pairs on oxygen; the lone pairs push the bonding pairs together, giving a bent geometry with a 104.5° bond angle. This shape creates a permanent dipole moment of 1.85 debye — the source of water's exceptional solvent properties and high boiling point. Ammonia (NH₃) has one lone pair and three bonding pairs, producing a trigonal pyramidal shape. Carbon dioxide (CO₂) has two double bonds and no lone pairs, yielding linear geometry and zero dipole moment despite having polar bonds.

Molecular shape governs biological activity: a drug molecule fits a receptor precisely because of its three-dimensional geometry. The lock-and-key model of enzyme catalysis — where substrate fits the enzyme active site — is a direct consequence of molecular shape arising from chemical bonding principles.