Why do atoms bond?
An atom has a nucleus in the middle and electrons in shells around it. The electrons in the outermost shell are called valence electrons. Only they take part in bonding.
Noble gases like neon and argon have 8 outer electrons. They almost never react. So chemists say: atoms bond to get 8 outer electrons, like a noble gas. This is the octet rule. Hydrogen and helium are small and are happy with 2 (a duplet).
The number of valence electrons comes from the group in the periodic table. Group 1 metals have 1, group 2 have 2, group 16 have 6, group 17 have 7.
Lewis dot symbols
A Lewis symbol shows the element symbol with dots for its valence electrons. Na has one dot. Cl has seven dots. Lewis structures help us see who gives, takes or shares.
Ionic bonding: giving and taking
Metals have few outer electrons and lose them easily. Non-metals have many and gain easily. When they meet, the metal transfers electrons to the non-metal.
- An atom that loses electrons becomes a positive ion (cation), e.g. Na → Na+ + e−.
- An atom that gains electrons becomes a negative ion (anion), e.g. Cl + e− → Cl−.
The strong attraction between opposite ions is the ionic bond. The ions pack into a giant 3D lattice. The formula shows the simplest ratio that makes the total charge zero: Mg2+ and Cl− give MgCl2; Ca2+ and O2− give CaO. Polyatomic ions like SO42− and NH4+ act as one unit: Na2SO4.
Properties of ionic compounds
- Hard, brittle crystals.
- High melting and boiling points (strong forces in every direction).
- Many dissolve in water.
- Conduct electricity only when molten or dissolved, because then ions can move.
Covalent bonding: sharing pairs
Two non-metals both want electrons, so neither gives. Instead they share a pair of electrons. One shared pair is a single bond (H–H). Two pairs make a double bond (O=O). Three pairs make a triple bond (N≡N).
Atoms joined this way form molecules, like H2O, CO2 and CH4. Pairs not used in bonds are lone pairs. Ball-and-stick models show molecules in 3D.
Polar and non-polar bonds
Electronegativity is how strongly an atom pulls shared electrons (Pauling scale: F = 4.0, O = 3.4, Cl = 3.2, H = 2.2, Na = 0.9). If two atoms pull equally (H–H), the bond is non-polar. If one pulls harder (H–Cl), electrons sit closer to it: a polar bond with small charges δ+ and δ−.
Rough guide by electronegativity difference (ΔEN): below about 0.4 non-polar covalent; about 0.4 to 1.8 polar covalent; above about 1.8 mostly ionic. It is a scale, not three boxes.
Properties of molecular (covalent) compounds
- Often gases, liquids or soft solids with low melting points: the forces between molecules are weak.
- Do not conduct electricity (no free ions or electrons).
- Exception: giant covalent structures like diamond and quartz are very hard with very high melting points.
Metallic bonding, and choosing materials
In a metal, each atom releases its outer electrons into a shared sea of delocalised electrons. The positive metal ions sit in neat rows. The attraction between the ions and the electron sea is the metallic bond.
- Good conductors of electricity and heat: free electrons move.
- Malleable and ductile: layers of ions can slide without breaking the bond.
- Shiny, and usually high melting points.
Why bonding matters in daily life
Bond type decides how we use a substance and how safe it is. Sodium chloride is safe in food but too much salt raises blood pressure. Covalent solvents in paints and nail polish evaporate easily (weak forces between molecules), so their vapour can be breathed in; use them in open air. Mercury is a metal that is liquid at room temperature and its vapour is poisonous, which is why old mercury thermometers are being replaced.
Key formulas and definitions
- Octet rule: atoms tend to gain, lose or share electrons to have 8 outer electrons (H: 2)
- Ionic formula: total positive charge = total negative charge (e.g. Al³⁺ + O²⁻ → Al₂O₃)
- ΔEN = |EN(A) − EN(B)|: < 0.4 non-polar, 0.4–1.8 polar, > 1.8 mostly ionic (rough guide)
- Single bond = 1 shared pair, double = 2 pairs, triple = 3 pairs
- Ionic = metal + non-metal; covalent = non-metals; metallic = metal atoms
Worked examples
1. Show how magnesium (2 outer electrons) and chlorine (7) form magnesium chloride. Write the formula.
Mg gives 1 electron to each of two Cl atoms: Mg → Mg²⁺ + 2e⁻; each Cl + e⁻ → Cl⁻. Charges: +2 and 2 × (−1) = 0. Formula MgCl₂.
2. Write the formula of the compound of aluminium (Al³⁺) and oxygen (O²⁻).
Find the lowest common multiple of 3 and 2: 6. Two Al³⁺ give +6, three O²⁻ give −6. Formula Al₂O₃.
3. Predict the bond type in (a) H–H, (b) H–O, (c) K–F. EN: H 2.2, O 3.4, K 0.8, F 4.0.
(a) ΔEN = 0 → non-polar covalent. (b) ΔEN = 1.2 → polar covalent. (c) ΔEN = 3.2 → ionic.
4. How many shared pairs are in N₂? Why?
N has 5 outer electrons and needs 3 more. Each N shares 3 electrons, so there are 3 shared pairs: a triple bond N≡N. Each N also keeps 1 lone pair.
5. Solid salt does not conduct electricity but salt water does. Explain.
In solid NaCl the ions are locked in the lattice and cannot move. In water the lattice breaks up and ions move freely, so they carry charge.
6. Why does copper bend into wire while salt crystals shatter?
In copper, layers of ions slide and the electron sea still holds them. In salt, a slide brings like charges side by side; they repel and the crystal cracks.
Common mistakes
- Saying ionic bonds are 'shared electrons'. Ionic = transfer; covalent = sharing.
- Thinking a solid ionic compound conducts. It conducts only when melted or dissolved, when ions can move.
- Believing covalent means weak. The covalent bond itself is strong; low melting points come from weak forces BETWEEN molecules. Diamond is covalent and very hard.
- Writing NaCl₂ or MgCl. Always balance charges: Na⁺ with Cl⁻ gives NaCl; Mg²⁺ needs two Cl⁻, so MgCl₂.