Revision notes: Chemical Bond

These notes work through the Chemical Bond chapter one content standard at a time, so you can revise each idea and check you have the definitions the marking scheme expects. Bonding is a chapter you learn by drawing: keep a pencil handy and sketch every dot-and-cross diagram as you read.

5.1 Why atoms form bonds, stability

Atoms bond because a full outer shell is a stable, low-energy arrangement. The noble gases in Group 18 already have a full outer shell, a duplet (2 electrons) for helium, an octet (8 electrons) for the others, so they are unreactive and exist as single atoms. Every other atom bonds to reach the same stable electron arrangement as the nearest noble gas.

Key idea. Atoms lose, gain or share electrons so that each atom ends up with a full outer shell. This is the octet rule (duplet for the smallest atoms such as hydrogen). It is the reason behind every bond in this chapter.

A metal atom has few outer electrons and reaches stability most easily by losing them; a non-metal atom has an almost-full outer shell and reaches stability by gaining or sharing electrons. That single difference decides which kind of bond forms.

5.2 The ionic bond, electron transfer

An ionic bond forms between a metal and a non-metal by the transfer of electrons. The metal atom loses its outer electrons to become a positively charged cation; the non-metal atom gains those electrons to become a negatively charged anion. The oppositely charged ions then attract by strong electrostatic forces, that attraction is the ionic bond.

Take sodium chloride. Sodium (electron arrangement 2.8.1) transfers one electron to chlorine (2.8.7). Sodium becomes Na⁺ (2.8) and chlorine becomes Cl⁻ (2.8.8); both now have a noble-gas arrangement, and Na⁺ and Cl⁻ attract to form NaCl.

Dot-and-cross rule. Use dots for one atom’s electrons and crosses for the other so the marker can see where each electron came from. Draw square brackets around each ion and write its charge outside the top-right corner: [Na]⁺ and [ :Cl: ]⁻.

Watch the ratio. Magnesium (2.8.2) loses two electrons; chlorine gains one each, so one Mg²⁺ needs two Cl⁻, the formula is MgCl₂. In magnesium oxide, Mg²⁺ and O²⁻ combine one-to-one as MgO.

5.3 The covalent bond, electron sharing

A covalent bond forms between non-metal atoms by sharing pairs of electrons. Each shared pair counts towards the octet of both atoms, so neither atom needs to lose or gain electrons outright.

• A single bond is one shared pair, as in H₂, Cl₂, HCl, H₂O, NH₃ and CH₄.
• A double bond is two shared pairs, as in O₂ and CO₂.
• A triple bond is three shared pairs, as in N₂.

In water, oxygen (2.6) shares one electron with each of two hydrogen atoms, giving oxygen an octet and each hydrogen a duplet; oxygen also keeps two lone pairs (unshared pairs). In methane, carbon (2.4) shares with four hydrogen atoms to complete its octet.

Dot-and-cross rule. Show only the outer electrons. Put the shared pair(s) in the overlap between the two atoms, and remember to draw the lone pairs too, leaving them off is a frequent error in O₂, N₂, H₂O and NH₃.

5.4 Dative bond and hydrogen bond

A dative (coordinate) bond is a covalent bond in which both shared electrons come from the same atom. Once formed it behaves exactly like an ordinary covalent bond. The standard examples are the ammonium ion NH₄⁺ (the lone pair on nitrogen in NH₃ bonds to an H⁺) and the hydronium ion H₃O⁺ (the lone pair on oxygen in H₂O bonds to an H⁺). Show a dative bond with an arrow pointing away from the donor atom.

A hydrogen bond is a weak attraction between molecules, not a bond within a molecule. It forms when hydrogen is bonded to a very electronegative atom, fluorine, oxygen or nitrogen, and is attracted to a lone pair on such an atom in a neighbouring molecule. Hydrogen bonding raises the boiling point of water and other molecules well above what their size alone would predict.

Do not confuse the two. A dative bond is a strong bond inside a molecule; a hydrogen bond is a weak attraction between separate molecules.

5.5 Properties from bonding

The whole point of the chapter is that structure and bonding explain physical properties. Learn these three comparisons.

• Melting and boiling point. Ionic compounds form a giant ionic lattice held by strong electrostatic forces throughout, so a large amount of heat is needed to separate the ions, high melting and boiling points. Simple covalent (molecular) compounds have strong bonds inside each molecule but only weak forces between molecules, so little heat separates the molecules, low melting and boiling points.
• Electrical conductivity. An ionic compound conducts electricity only when molten or in aqueous solution, because the ions are then free to move; as a solid the ions are locked in place and it does not conduct. Simple covalent compounds have no free-moving ions or electrons, so they do not conduct in any state (graphite is the well-known exception).
• Solubility. Ionic compounds generally dissolve in water but not in organic solvents; many simple covalent compounds dissolve in organic solvents but not in water. “Like dissolves like” is a useful reminder.

Exam habit. When you explain a property, always name the structure and the force. “High melting point because strong electrostatic forces of attraction between oppositely charged ions in the giant lattice need a lot of energy to overcome.” The reason carries the marks.

How to use these notes

Revise one standard at a time and redraw every diagram from memory. For 5.1, state the octet rule and why noble gases are unreactive. For 5.2 and 5.3, practise dot-and-cross diagrams for NaCl, MgO, MgCl₂, H₂O, NH₃, CH₄ and CO₂ until they are automatic. For 5.4, be able to explain where the shared pair in a dative bond comes from. For 5.5, rehearse the three property comparisons with the force named each time. A one-to-one teacher can check that your diagrams show charges, brackets and lone pairs correctly, the small details that separate a part-mark from full marks across SPM Chemistry.