Why atoms bond
Atoms are most stable when their outer electron shell is full: usually eight electrons (an octet), or two for hydrogen. The noble gases already have full shells, which is why they barely react. Other atoms reach a full shell by sharing electrons (covalent bonding) or by transferring them (ionic bonding). Bonding lowers the energy of the atoms, so energy is released when bonds form.
Covalent bonds
A covalent bond is a pair of electrons shared between two atoms, usually non-metals. Both nuclei attract the shared pair, holding the atoms together. Atoms can share one, two or three pairs, formingsingle, double or triple bonds; more shared pairs make shorter, stronger bonds. Outer electron pairs that are not shared are called lone pairs.
Ionic bonds
When a metal reacts with a non-metal, the metal atom usually transfers its outer electrons to the non-metal. The metal becomes a positive ion (such as Na⁺ or Mg²⁺) and the non-metal a negative ion (Cl⁻ or O²⁻). These oppositely charged ions attract strongly in every direction, building a giant ionic lattice. There are no separate molecules: the formula NaCl just gives the ratio of ions.
Electronegativity and polarity
Electronegativity measures how strongly an atom attracts shared electrons. On Linus Pauling’s scale, fluorine is highest (3.98) and the alkali metals lowest. The difference between two atoms predicts the bond type:
- Less than about 0.4: non-polar covalent (electrons shared evenly), as in H₂ or C–H bonds.
- About 0.4 to 1.8: polar covalent: the more electronegative atom gets a partial negative charge (δ−), as in O–H.
- More than about 1.8: ionic: the electron is effectively transferred, as in NaCl.
These cut-offs are a guide; bonding is really a spectrum. Whether a whole molecule is polar also depends on its shape: the polar bonds in CO₂ cancel because the molecule is straight.
Molecular shapes (VSEPR)
Electron pairs around a central atom repel each other and spread out as far as possible. This isvalence shell electron pair repulsion (VSEPR) theory. Four pairs point to the corners of a tetrahedron (109.5°), as in methane. Lone pairs repel more strongly than bonding pairs, squeezing the bond angles: 107° in ammonia (one lone pair) and 104.5° in water (two lone pairs). Two pairs or double bonds on opposite sides give a linear molecule, like CO₂.
Structure and properties
| Structure | Example | Melting point | Conducts electricity? |
|---|---|---|---|
| Simple molecular | Methane, water | Low (weak forces between molecules) | No |
| Giant ionic | Sodium chloride | High | Only when melted or dissolved |
| Giant covalent | Diamond | Very high | No (graphite is an exception) |
Key terms
- Ion
- An atom or group of atoms that has lost or gained electrons, giving it a charge.
- Lone pair
- A pair of outer-shell electrons not involved in bonding.
- Dipole
- A separation of charge, with a slightly positive end and a slightly negative end.
Common misconceptions
- “Boiling water breaks the O–H bonds.” It only overcomes the forces between molecules; the molecules stay intact.
- “Salt is made of NaCl molecules.” It is a continuous lattice of ions.
- “Ionic and covalent are completely separate.” Bonds range smoothly from equal sharing to complete transfer.
Key takeaways
- Covalent bonds share electron pairs; ionic bonds transfer electrons to form oppositely charged ions.
- Electronegativity differences predict whether a bond is non-polar, polar or ionic.
- Repulsion between electron pairs, especially lone pairs, sets a molecule’s shape.
- Structure explains properties: small molecules melt easily; giant lattices and networks do not.
Model note: atoms are drawn as spheres and bonds as sticks; real electrons form fuzzy clouds. Lone-pair lobes are shown approximately. Only a small piece of each giant structure is drawn.