Chemical Bonding: Ionic, Covalent, Metallic and Hydrogen

Chemical bonding is the force that holds atoms together in molecules and materials. Atoms bond because doing so usually gives them lower, more stable energy — most often by achieving a full outer shell of electrons. The way electrons are shared or transferred creates four main bond types: ionic, covalent, metallic and hydrogen bonding.

Bonding matters because it decides the properties of everything around you. Why salt dissolves in water but oil does not, why copper conducts electricity, why diamond is hard and wax is soft — all of it traces back to which bonds hold the atoms together.

Key Concepts

• Octet rule: atoms gain, lose or share electrons to reach eight in their outer shell, like the noble gases.
• Electronegativity: an atom's pull on shared electrons. The difference predicts bond type: below ~0.4 nonpolar covalent, 0.4–1.7 polar covalent, above ~1.7 ionic.
• Ionic bond: complete electron transfer from metal to non-metal, forming oppositely charged ions that attract. Example: NaCl.
• Covalent bond: sharing of electron pairs. Equal sharing gives nonpolar bonds (H₂); unequal sharing gives polar bonds with partial charges (H₂O).
• Metallic bond: a "sea" of delocalised electrons among positive metal ions — the reason for conductivity and malleability.
• Hydrogen bond: a weak attraction between H bonded to N, O or F and a lone pair on another N, O or F. Weak alone, powerful in numbers — it gives water its unusual properties.

Formulas

• Bond order = (bonding electrons − antibonding electrons) / 2. Dimensionless: 1 = single, 2 = double, 3 = triple bond. Higher bond order means a shorter, stronger bond.
• Balanced equations like 2H₂ + O₂ → 2H₂O conserve atoms; adjust coefficients, never subscripts.

Worked Example

Find the bond order of O₂, which has 10 bonding and 6 antibonding electrons.

1. Bond order = (10 − 6) / 2 = 4 / 2 = 2.
2. A bond order of 2 means O₂ has a double bond, O=O — matching its measured short bond length and high bond energy.

The arithmetic confirms experiment: oxygen's atoms are held by two shared pairs, making O₂ far less reactive than a single bond would allow.

Common Misconceptions

• "A hydrogen bond is a strong chemical bond." It is a weak intermolecular attraction — roughly one-twentieth the strength of a covalent bond. Water's properties come from billions acting together.
• "Electrons are always shared equally in covalent bonds." Only between identical atoms; usually the more electronegative atom hogs them, creating the partial charges that make water polar.
• "Ionic and covalent are totally separate." Bonding is a continuum: as electronegativity difference grows, a bond slides from covalent through polar to ionic.

Real-World Applications

• Materials: metallic bonding gives conductivity; covalent networks make diamond and silicon hard.
• Biology: hydrogen bonds hold DNA strands together and fold proteins.
• Drug design: medicines are shaped to form hydrogen bonds with target proteins.

Atoms bond to reach lower energy, and the electron arrangement decides the type: transfer makes ionic bonds, sharing makes covalent, delocalisation makes metallic, and weak attractions to hydrogen make hydrogen bonds. Bond order then tells you how strong and short the bond is.