How are atoms held together, and how does the type of bonding explain a material's properties?
Ionic, covalent, dative covalent and metallic bonding, electronegativity and bond polarity, electron-pair repulsion and the shapes of simple molecules and ions, and the link between structure and physical properties.
A CCEA A-Level Chemistry answer on ionic, covalent, dative covalent and metallic bonding, electronegativity and bond polarity, electron-pair repulsion theory and the shapes of simple molecules and ions, and how structure determines physical properties.
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What this dot point is asking
CCEA wants you to describe ionic, covalent, dative covalent and metallic bonding, explain electronegativity and bond polarity, use electron-pair repulsion theory to predict and name the shapes and bond angles of simple molecules and ions, and link structure to physical properties such as melting point, conductivity and solubility.
Types of bonding
Electronegativity and polarity
A molecule with polar bonds can still be non-polar overall if the dipoles cancel by symmetry, as in and . Electronegativity follows clear periodic trends: it increases across a period as nuclear charge rises and atomic radius falls, and decreases down a group as the bonding electrons sit further from the nucleus and are more shielded. Fluorine is the most electronegative element. When the difference in electronegativity becomes very large (a metal bonded to a non-metal), the bond is essentially ionic; when it is zero (two identical atoms, as in ), the bond is pure covalent. Most bonds lie somewhere in between, described as polar covalent.
Shapes of molecules and ions
Electron-pair repulsion theory states that electron pairs around a central atom repel each other and arrange themselves as far apart as possible. Lone pairs repel more strongly than bonding pairs, reducing bond angles by about each.
Structure and properties
Giant ionic lattices have high melting points and conduct when molten or dissolved. Giant covalent structures (diamond, graphite, silicon dioxide) have very high melting points; graphite conducts because of delocalised electrons. Simple molecular substances have low melting points because only weak intermolecular forces are broken on melting. Metals conduct and are malleable because of mobile delocalised electrons.
Why each structure behaves as it does
In a giant ionic lattice such as sodium chloride, every ion is held by strong electrostatic attraction to its oppositely charged neighbours in all directions, so a large amount of energy is needed to melt it. It does not conduct when solid because the ions are locked in place, but it conducts when molten or in solution because the ions become free to move and carry charge. Ionic solids tend to be brittle: a layer of ions shifted by one position brings like charges next to each other, which repel and split the crystal.
In a giant covalent structure such as diamond, every carbon is covalently bonded to four others throughout the crystal, so the melting point is extremely high and the solid does not conduct (no free electrons). Graphite is the exception: each carbon bonds to only three others, leaving one delocalised electron per atom free to move along the layers, which makes graphite a good conductor and a soft lubricant because the layers slide.
In metals, the delocalised electrons are free to move, carrying both charge (conduction) and heat. The layers of positive ions can slide over one another without breaking the bonding, which is why metals are malleable and ductile.
Examples in context
Example 1. Choosing a material for an electrode. Graphite electrodes are used in electrolysis even though graphite is a non-metal. The reason lies in the structure: graphite has delocalised electrons free to move between its layers, so it conducts electricity, while being chemically inert and high melting. A CCEA candidate is expected to link this directly to graphite's bonding (three covalent bonds per carbon plus one delocalised electron) rather than just stating that it conducts.
Example 2. Predicting whether a solvent will dissolve a salt. Water dissolves ionic compounds such as sodium chloride because water is a polar molecule (its bent shape means the bond dipoles do not cancel) and can surround and stabilise the separated ions. Non-polar solvents such as hexane cannot, because they offer no charge to interact with the ions. The argument runs straight through this dot point: bond polarity to molecular shape to overall polarity to solubility, which is exactly the chain of reasoning CCEA questions test.
Try this
Q1. State the shape and bond angle of an ammonia molecule. [2 marks]
- Cue. Pyramidal, with a bond angle of about .
Q2. Explain why magnesium chloride conducts electricity when molten but not when solid. [2 marks]
- Cue. When molten the ions are free to move and carry charge; in the solid the ions are fixed in the lattice.
Exam-style practice questions
Practice questions written in the style of CCEA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
CCEA 20196 marksSulfur dioxide () and the sulfite ion () both contain a central sulfur atom. State the shape and bond angle of each species and explain, using electron-pair repulsion theory, why their bond angles differ.Show worked answer β
Markers want the electron-pair count for each species, the resulting shape and angle, and a comparison.
Sulfur dioxide has two bonding regions and one lone pair on the central sulfur, so three regions of electron density. This gives a bent (V-shaped) molecule. With one lone pair, the bond angle is about , slightly below the of a pure trigonal arrangement.
The sulfite ion has three bonding regions and one lone pair, so four regions of electron density. This gives a pyramidal ion with a bond angle of about .
The key explanation: lone pairs repel more strongly than bonding pairs, and the sulfite ion has four electron regions arranged tetrahedrally compressed by one lone pair, whereas sulfur dioxide has three regions in a plane compressed by one lone pair. More electron regions and the extra bonding pair push the sulfite angle well below the dioxide angle.
A common error is to ignore the lone pair on sulfur dioxide and call it linear like carbon dioxide.
CCEA 20214 marksExplain why magnesium has a higher melting point than sodium, and why both are much higher melting than the noble gas argon.Show worked answer β
This tests the link between metallic bonding strength and physical properties.
Magnesium forms ions and releases two delocalised electrons per atom, whereas sodium forms and releases one. Magnesium therefore has a greater charge on its ions and more delocalised electrons, so the electrostatic attraction between the ions and the sea of electrons is stronger. Its ions are also smaller, bringing the charges closer. Both factors mean more energy is needed to break the metallic bonding, so magnesium melts higher.
Argon is monatomic with no metallic bonding at all; the only forces between its atoms are very weak van der Waals (London) forces, which need very little energy to overcome, so argon has an extremely low melting point.
Markers reward the charge and electron-number comparison for the metals and the contrast with weak van der Waals forces in argon.
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Sources & how we know this
- CCEA GCE Chemistry specification β CCEA (2016)