Chapter 3: Chemical Bonding

Ionic bonding occurs when there is a large electronegativity difference between atoms, typically a metal and a non-metal. The metal loses valence electrons to form a cation, and the non-metal gains electrons to form an anion. The resulting ions attract each other by strong electrostatic (Coulombic) forces, producing an ionic lattice. Sodium chloride illustrates this: Na transfers one electron to Cl, giving Na⁺Cl⁻ in a crystal lattice.

Covalent bonding occurs when atoms share electron pairs so that each atom approaches a stable electron configuration. Sharing may be equal (non-polar covalent) or unequal (polar covalent) depending on electronegativity. Molecules such as H₂, Cl₂, H₂O and CH₄ are held by covalent bonds. In solids, covalent bonding may form discrete molecules or giant networks as in diamond and silicon dioxide.

Ionic compounds generally have high melting and boiling points because lattice energy is large. They conduct electricity when molten or dissolved because ions are free to move, but not in the solid state. Many ionic solids are brittle and dissolve in polar solvents such as water.

Covalent molecular substances often have lower melting points because intermolecular forces are weaker than ionic lattice forces. They do not conduct electricity in the pure molecular state. Polar covalent compounds may dissolve in polar solvents; non-polar ones dissolve in non-polar solvents. Giant covalent solids are hard and have very high melting points.

In practice, bonding is a continuum: few compounds are purely ionic or purely covalent. Fajans’ rules help judge polarisation of ions and partial covalent character in ionic compounds. Electronegativity difference, polarisability, and structure together decide observed properties.

Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron pairs around a central atom arrange themselves to minimise repulsion. Both bonding pairs and lone pairs occupy space, but lone-pair–lone-pair repulsion is greater than lone-pair–bonding-pair repulsion, which is greater than bonding-pair–bonding-pair repulsion. The arrangement of electron pairs gives the electron-pair geometry; the molecular shape considers only atom positions.

In BF₃, boron has three bonding pairs and no lone pair. The electron pairs form a trigonal planar arrangement with bond angles of 120°. The molecular shape is therefore trigonal planar.

In CH₄, carbon has four bonding pairs and no lone pair. The preferred arrangement is tetrahedral with bond angles of about 109.5°. Methane is tetrahedral.

In NH₃, nitrogen has three bonding pairs and one lone pair. The electron-pair geometry is tetrahedral, but the molecular shape is trigonal pyramidal. The H–N–H angle is slightly less than 109.5° because the lone pair compresses the bonding pairs.

In H₂O, oxygen has two bonding pairs and two lone pairs. Electron-pair geometry remains tetrahedral, but the molecular shape is bent (V-shaped). The H–O–H angle is about 104.5° due to strong lone-pair repulsion.

VSEPR successfully predicts many main-group molecular shapes from Lewis structures. It does not by itself give bond lengths or explain magnetism; those need other models such as MOT.

Molecular orbital (MO) theory constructs molecular orbitals by linear combination of atomic orbitals. Electrons occupy MOs of increasing energy, following Aufbau, Pauli, and Hund rules. Bonding MOs lower energy; antibonding MOs raise energy. Bond order equals number of bonding electrons − number of antibonding electrons2.

For N₂ (14 valence electrons), the relevant MOs fill so that all electrons are paired. The bond order of N₂ is 3, matching the strong N≡N triple bond. Because there are no unpaired electrons, N₂ is diamagnetic and is weakly repelled by a magnetic field.

For O₂ (16 valence electrons), two additional electrons enter the π2p antibonding orbitals. Hund’s rule places one electron in each of the two degenerate π orbitals with parallel spins. Thus O₂ has two unpaired electrons and is paramagnetic, attracted into a magnetic field. The bond order of O₂ is 2, consistent with a double bond.

Simple Lewis structures for O₂ show paired electrons and wrongly predict diamagnetism. MOT corrects this and also explains the relative bond orders of O₂⁺, O₂, O₂⁻ and O₂²⁻. Experimentally, liquid oxygen is attracted to a magnet, confirming paramagnetism.

Therefore MOT is essential for molecules where Lewis diagrams fail, especially for magnetism and bond-order trends in diatomic second-period molecules.

In an ideal ionic compound, ions are spherical and held only by electrostatic attraction. In reality, a cation can distort the electron cloud of an anion; this distortion is polarisation. If polarisation is strong, electron density shifts toward the cation and the bond gains covalent character.

Fajans’ rules summarise when polarisation is large. Small, highly charged cations have high polarising power (e.g. Al³⁺, Be²⁺). Large, highly charged anions are easily polarisable (e.g. I⁻, S²⁻). Cations with a pseudo-noble-gas configuration (e.g. Cu⁺, Ag⁺, Zn²⁺) polarise anions more than similarly sized noble-gas-configuration cations.

Consequently, LiCl has more covalent character than NaCl, and AlCl₃ is markedly covalent compared with NaCl. Compounds of large anions such as iodides often show more covalent behaviour than fluorides of the same cation. Polarisation also affects colour, solubility, melting point, and conductivity trends.

Electronegativity difference alone does not decide all properties; polarisation explains many exceptions. For example, AgCl is less soluble and more covalent in character than NaCl despite both being formally ionic. Fajans’ approach links ionic size and charge to the ionic–covalent continuum.

Students should treat ionic and covalent labels as extremes on a continuous scale rather than absolute categories.

Hydrogen bonding is a strong directional intermolecular attraction that occurs when hydrogen is bonded to a highly electronegative atom such as F, O, or N. The H atom becomes electron-deficient and is attracted to a lone pair on a nearby electronegative atom. Hydrogen bonds are weaker than covalent bonds but stronger than typical van der Waals forces.

In water, each molecule can form up to four hydrogen bonds through two O–H bonds and two lone pairs on oxygen. In liquid water these bonds continually break and reform, giving a dynamic, partially ordered network. Hydrogen bonding explains water’s anomalously high boiling point, high specific heat, and excellent solvent properties for polar and ionic solutes.

When water freezes, the hydrogen-bonded network locks into an open hexagonal lattice. Each oxygen is tetrahedrally surrounded, and empty space remains within the structure. Because of this open structure, ice occupies a larger volume than the same mass of liquid water near 0 °C. Therefore ice is less dense than liquid water and floats on water.

On melting, the open lattice partially collapses and molecules pack more closely, so density increases from ice to liquid water. This anomaly is biologically and environmentally important: lakes freeze from the top, insulating aquatic life. HF and NH₃ also show hydrogen bonding, but water’s ability to form four bonds per molecule makes its effects especially pronounced.

Thus hydrogen bonding controls both the structure of ice and many unique properties of water.