Worked examples: Chemical Bond

Work through each example by covering the answer, attempting it yourself, then checking your method against the solution. Every example flags the slip that most often costs marks. These are original questions in SPM style, not past-year papers.

Example 1, Ionic bonding in sodium chloride

Question. Sodium has electron arrangement 2.8.1 and chlorine 2.8.7. Describe, with reference to electron transfer, how sodium chloride forms.

Solution. The sodium atom transfers its one outer electron to the chlorine atom. Sodium becomes a sodium ion Na⁺ with arrangement 2.8; chlorine becomes a chloride ion Cl⁻ with arrangement 2.8.8. Both ions now have a stable noble-gas arrangement. The oppositely charged Na⁺ and Cl⁻ are held together by strong electrostatic forces of attraction, the ionic bond, giving NaCl.

Common slip: saying atoms “share” electrons in an ionic bond. Ionic bonding is transfer, not sharing.

Example 2, Getting the ion ratio right (magnesium chloride)

Question. Magnesium is 2.8.2 and chlorine is 2.8.7. Explain why the formula of magnesium chloride is MgCl₂.

Solution. A magnesium atom loses two electrons to become Mg²⁺ (2.8). Each chlorine atom can gain only one electron to become Cl⁻ (2.8.8). To place both of magnesium’s electrons, two chlorine atoms are needed. So one Mg²⁺ combines with two Cl⁻, and the overall charge balances to zero: the formula is MgCl₂.

Common slip: writing MgCl because the charges (2+ and 1−) were not balanced. Always check the total positive charge equals the total negative charge.

Example 3, Dot-and-cross for water

Question. Oxygen is 2.6 and hydrogen is 2.1… (that is, hydrogen has 1 electron). Draw and describe the dot-and-cross diagram of a water molecule, H₂O.

Solution. Oxygen shares one of its outer electrons with each hydrogen atom, and each hydrogen shares its single electron with oxygen. This makes two single covalent bonds (two shared pairs). Oxygen now has eight electrons around it (an octet) and each hydrogen has two (a duplet). Oxygen also carries two lone pairs that are not shared.

Common slip: forgetting the two lone pairs on oxygen, or drawing eight crosses around oxygen instead of showing which electrons are shared.

Example 4, A double bond (carbon dioxide)

Question. Explain the bonding in carbon dioxide, CO₂, and state how many electrons are shared.

Solution. Carbon (2.4) shares two electron pairs with each oxygen atom, forming a double covalent bond on each side (O=C=O). Carbon reaches an octet by sharing four pairs in total; each oxygen reaches an octet through its two shared pairs plus its two lone pairs. Four electrons are shared in each C=O bond.

Common slip: drawing single bonds in CO₂. Carbon has only four outer electrons, so single bonds would leave it short of an octet, double bonds are required.

Example 5, A dative bond (ammonium ion)

Question. Explain how the fourth N–H bond in the ammonium ion, NH₄⁺, is different from the other three.

Solution. In ammonia, NH₃, nitrogen has one lone pair. When NH₃ reacts with a hydrogen ion H⁺ (which has no electrons), the nitrogen lone pair provides both electrons for the new bond. This is a dative (coordinate) bond. Once formed, it is identical in strength to the other three N–H bonds; the difference is only in the origin of the shared pair.

Common slip: saying the H⁺ brings an electron to share. H⁺ has no electron; both electrons come from nitrogen.

Example 6, Hydrogen bonding and boiling point

Question. Water has an unexpectedly high boiling point for such a small molecule. Explain why in terms of bonding.

Solution. In water, hydrogen is bonded to the highly electronegative oxygen, so each molecule can form hydrogen bonds with neighbouring water molecules. These extra attractions between molecules must be overcome before the water boils, so more energy, and a higher temperature, is needed than the small molecular size alone would suggest.

Common slip: claiming the hydrogen bonds inside the molecule must be broken. Boiling separates molecules, so it is the hydrogen bonds between molecules that matter.

Example 7, Explaining melting point from bonding

Question. Sodium chloride has a high melting point while iodine has a low one. Explain the difference in terms of structure and bonding.

Solution. Sodium chloride is a giant ionic lattice: strong electrostatic forces act between the oppositely charged ions throughout the solid, and a large amount of energy is needed to overcome them, so the melting point is high. Iodine is a simple molecular solid: the covalent bonds within each I₂ molecule are strong, but the forces between molecules are weak, so little energy separates them and the melting point is low.

Common slip: saying that covalent bonds are broken when iodine melts. Melting overcomes the weak forces between molecules, not the strong bonds within them.

Example 8, Explaining electrical conductivity

Question. Explain why solid lead(II) bromide does not conduct electricity but molten lead(II) bromide does.

Solution. Lead(II) bromide is an ionic compound. As a solid, the ions are held in fixed positions in the lattice and cannot move, so there are no mobile charge carriers and it does not conduct. When molten, the lattice breaks down and the ions become free to move, carrying charge through the liquid, so it conducts electricity.

Common slip: saying electrons carry the charge. In an ionic compound it is the mobile ions, not electrons, that conduct.

Using these examples

Every answer here does one of three things: it names the electron transfer or sharing, it justifies the ion ratio or bond type, or it links a property to structure and the force involved, the habits this chapter rewards. Once the reasoning feels automatic, move on to the practice questions and mark yourself the same way. A one-to-one teacher can check that your dot-and-cross diagrams show charges, brackets and lone pairs, which is exactly where marks are gained or lost across the SPM Chemistry written papers.

Example 9, Explaining solubility

Question. Sodium chloride dissolves readily in water but naphthalene (a simple molecular solid) does not. Explain in terms of bonding.

Solution. Water is a polar solvent. Sodium chloride is ionic, and its ions are attracted to the polar water molecules, so the lattice breaks up and the ions spread through the water, it dissolves. Naphthalene is a non-polar molecular solid held by weak forces between molecules; water cannot pull these molecules apart, so it does not dissolve. Naphthalene would, however, dissolve in a non-polar organic solvent, following the rule “like dissolves like”.

Common slip: saying an ionic compound dissolves in every solvent. Ionic solids dissolve in water but generally not in non-polar organic solvents.