Chemical Bonding

Atoms combine in order to become more stable, and stability for an atom means having a completely filled outermost shell, just like the noble gases.

• Atoms bond to lower their energy and become more stable; a lower-energy arrangement is a more stable one.
• The noble gases (He, Ne, Ar, etc.) are chemically inert because their outermost shells are already complete.
• Other atoms try to achieve the stable electronic configuration of the nearest noble gas by gaining, losing or sharing electrons.

The octet rule sums up the goal of bonding: every atom "wants" eight electrons in its outermost shell, the same count that makes noble gases so stable.

• Atoms gain, lose or share electrons to complete their octet and reach a noble-gas configuration.
• For the lightest atoms like hydrogen, the target is a duplet of 2 electrons (the configuration of helium), not an octet.
• Examples: Na loses 1 electron and Cl gains 1 electron — both then have a full octet.

Valency is the simple number that tells us how many bonds an atom can form, and it follows directly from its outermost electrons.

• Valency depends on the number of valence (outermost) electrons.
• An atom with 1, 2 or 3 valence electrons usually loses them, so valency = number of valence electrons.
• An atom with 5, 6 or 7 valence electrons usually gains electrons, so valency = 8 − valence electrons.
• Examples: Na (1 valence electron) has valency 1; O (6 valence electrons) has valency 2; Al has valency 3.

Chemical bonds come in a few clear types, each formed by a different way of dealing with electrons.

An ionic bond is formed when one atom hands over electrons to another, turning both into oppositely charged ions that then attract each other.

• It forms between a metal and a non-metal — the metal loses electrons, the non-metal gains them.
• The atom that loses electrons becomes a positive ion (cation); the atom that gains electrons becomes a negative ion (anion).
• The bond is the electrostatic attraction between these oppositely charged ions.
• Example — NaCl: sodium (Na) gives its 1 valence electron to chlorine (Cl), forming Na⁺ and Cl⁻, which attract to form common salt.
• Other examples: MgCl₂, CaO, KCl, MgO are all ionic compounds.

Because ionic compounds are built from a rigid lattice of charged ions, they share a distinctive set of physical properties.

• They are usually solids at room temperature, hard and brittle.
• They conduct electricity in the molten or aqueous (dissolved) state, where ions are free to move — but NOT in the solid state.
• They are generally soluble in water and insoluble in organic solvents like kerosene or petrol.
• They do not show the property of isomerism and react quickly in solution (ionic reactions are fast).

A covalent bond is formed when two atoms share electrons instead of transferring them, each contributing to a shared pair so that both complete their octet.

• It forms between two non-metal atoms (which both tend to gain electrons, so neither will simply give them away).
• Each shared pair of electrons counts towards the octet of both atoms.
• A covalent bond may involve sharing of one, two or three pairs of electrons.

Covalent bonds are classified by how many electron pairs are shared, and more shared pairs mean a stronger, shorter bond.

• H₂ — two hydrogen atoms share 1 pair (a single bond), each reaching the helium duplet.
• O₂ — two oxygen atoms share 2 pairs (a double bond).
• N₂ — two nitrogen atoms share 3 pairs (a triple bond), one of the strongest bonds known.
• H₂O — oxygen forms a single bond with each of two hydrogen atoms.
• CO₂ — carbon forms a double bond with each of two oxygen atoms (O=C=O).

Covalent compounds exist as discrete molecules held together only weakly, which gives them properties almost opposite to those of ionic compounds.

• They are often gases, liquids or soft solids at room temperature.
• They are poor conductors of electricity (bad conductors) as they have no free ions or electrons.
• They are mostly insoluble in water but generally soluble in organic solvents.
• Covalent reactions are slow and the compounds often show isomerism.
• Exception: diamond is a covalent solid that is extremely hard with a very high melting point — a giant covalent network of carbon atoms.

A coordinate bond is a special covalent bond in which the shared pair of electrons is supplied entirely by just one of the two atoms.

• The atom that donates the electron pair is the donor; the atom that accepts it is the acceptor.
• Once formed, it is identical to an ordinary covalent bond and is shown by an arrow (→) pointing from donor to acceptor.
• Examples: the ammonium ion NH₄⁺ (the N lone pair bonds with H⁺) and the hydronium ion H₃O⁺.

In a metal, the atoms release their outer electrons into a common pool, and the bond is the attraction between these mobile electrons and the fixed positive ions.

• The valence electrons are not fixed to any one atom — they move freely throughout the metal.
• The free electrons explain why metals conduct heat and electricity well.
• The sliding layers of ions make metals malleable and ductile (they can be hammered into sheets and drawn into wires).
• The strong attraction also gives metals their metallic lustre and high melting points.

A side-by-side comparison makes the contrast between the two main bond types easy to remember for the exam.