Chemical bonding

Why this matters

Table salt melts at 801 °C and dissolves in water; candle wax melts in your hand and does not. Steel conducts electricity; plastic insulates it. These differences come from how atoms are held together. Predicting the bond type from two element symbols lets you guess whether a substance will dissolve, conduct, melt easily or stay solid, which is everyday reasoning in labs, materials work and engineering.

The core idea: who holds the electrons?

A bond forms because atoms are more stable when they reach a full outer shell of electrons. They get there by transferring electrons or sharing them. Which one happens depends on how unequally the two atoms pull on electrons, measured by electronegativity (Pauling scale, from about 0.7 to 3.98).

A rule of thumb with electronegativity difference (ΔEN)

ΔEN < 0.4         nonpolar covalent
0.4 ≤ ΔEN < ~1.8  polar covalent
ΔEN ≥ ~1.8        mostly ionic

The cut-offs are approximate: bonding is a spectrum, and textbooks place the ionic boundary anywhere from about 1.7 to 2.0. The metal + nonmetal shortcut usually gives the same answer.

Worked examples

Electronegativity values: Na 0.93, Mg 1.31, H 2.20, C 2.55, N 3.04, O 3.44, Cl 3.16, F 3.98.

NaCl (sodium chloride). ΔEN = 3.16 − 0.93 = 2.23, and it is a metal with a nonmetal: ionic. Sodium gives its single valence electron to chlorine, making Na⁺ and Cl⁻. The ions stack into a crystal lattice held by strong attractions in every direction, which is why salt is hard, brittle and has a high melting point. Molten or dissolved, the ions can move, so it conducts electricity.

MgO (magnesium oxide). ΔEN = 3.44 − 1.31 = 2.13: ionic. Mg²⁺ and O²⁻ carry double charges, so the attraction is even stronger; MgO melts at about 2,850 °C and is used to line furnaces.

H₂O (water). O–H: ΔEN = 3.44 − 2.20 = 1.24: polar covalent. Oxygen pulls the shared electrons toward itself, so oxygen is δ− and each hydrogen δ+. Because the molecule is bent, those partial charges do not cancel, and water is a polar molecule. That polarity is why water dissolves salt and sugar so well.

CH₄ (methane). C–H: ΔEN = 2.55 − 2.20 = 0.35: nonpolar covalent. Methane, like wax and oils, is mostly C–H bonds, which is why it does not mix with water.

Cl₂ (chlorine gas). Two identical atoms, ΔEN = 0: nonpolar covalent.

CO₂ (carbon dioxide). Each C=O bond is polar (ΔEN = 0.89), but the molecule is straight (O=C=O), so the two pulls cancel. CO₂ has polar bonds but is a nonpolar molecule. Shape matters.

Single, double and triple bonds

Atoms can share one, two or three pairs of electrons:

• single: H–H, C–H
• double: O=O, C=O
• triple: N≡N (nitrogen gas, very hard to break, which is why N₂ is unreactive)

More shared pairs means a shorter, stronger bond.

Properties follow from bond type

"Like dissolves like" is the practical summary: polar and ionic substances mix with water; nonpolar substances mix with oils and other nonpolar solvents.

Common mistakes

• Saying ionic compounds are made of molecules. Solid NaCl is a lattice of ions; "NaCl" is a ratio, not a molecule.
• Treating polar bonds as a polar molecule. CO₂ and CCl₄ have polar bonds but symmetric shapes, so they are nonpolar.
• Using the ΔEN cut-off as a hard law. It is a guide. Hydrogen fluoride (ΔEN 1.78) is covalent even though it sits near the boundary.
• Forgetting metallic bonding. Two metals do not form an ionic bond; they form an alloy held by metallic bonding.

Quick recap

• Bonds form by transferring (ionic) or sharing (covalent) electrons.
• Electronegativity difference predicts the type; metal + nonmetal is usually ionic.
• Polar bonds in a symmetric molecule can cancel out.
• Bond type explains melting point, conductivity and solubility.

Try it

Open the Do tab and take the quiz: you will classify bonds from element pairs and predict properties from bond type.