Ionic Bonding and Ionic Compounds
In 1913, William and Lawrence Bragg used X-rays to map the NaCl crystal, revealing a perfect lattice of Na⁺ and Cl⁻ ions separated by just 0.28 nanometres.
Printable Worksheets
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Q1 · Think about solid table salt versus salt dissolved in water, which one do you think would conduct electricity, and why might there be a difference?
Q2 · Why do you think opposite charges attract each other, and how might that idea explain why ionic compounds like salt are held together so strongly?
● Know
- How ionic bonds form through electron transfer
- The structure of an ionic lattice
- The properties of ionic compounds (high mp, brittle, conducts when dissolved/melted)
● Understand
- Why ionic compounds have high melting points
- Why solid NaCl does not conduct electricity but dissolved NaCl does
- Why ionic compounds are brittle
● Can do
- Draw and explain the formation of an ionic bond
- Predict the properties of an ionic compound from its structure
- Explain the conductivity misconception for solid ionic compounds
Touch a freshly cut slice of ice and a copper pan that have both been sitting in a 0 °C room: the copper feels colder even at the same temperature, because copper's electrons move freely and carry heat away from your fingers instantly. A covalent bond forms when two non-metal atoms share a pair of electrons, with each atom contributing one electron to the shared pair. Unlike ionic bonding (which transfers electrons), covalent bonding shares them. Each atom still achieves a noble gas electron configuration by counting both its own electrons and the shared pair as its own. A single bond involves one shared pair; a double bond involves two shared pairs; a triple bond involves three shared pairs. Dot-cross diagrams show the electrons from each atom with different symbols (· vs ×) to track where each electron came from.
Water (H₂O) is a classic example: oxygen has 6 valence electrons and needs 2 more. Each hydrogen has 1 electron and needs 1 more. Oxygen forms single covalent bonds with two hydrogen atoms, each bond provides one extra electron to both the O and H. Oxygen achieves 8 valence electrons (neon configuration); each hydrogen achieves 2 (helium configuration). The resulting molecule H₂O is stable, with the electrons remaining in the region between the bonded atoms, the strongest attractive position between two positively charged nuclei.
Nitrogen gas (N₂): each nitrogen has 5 valence electrons and needs 3 more to reach 8. Two nitrogen atoms share 3 pairs of electrons, a triple bond. This triple bond is extraordinarily strong (945 kJ/mol to break), which is why nitrogen gas is so unreactive and makes up 78% of Australia's atmosphere without reacting with oxygen.
Liquid nitrogen (N₂, covalently bonded) is produced in large quantities at BOC Gas's Australian plants for food freezing, medical procedures, and laboratory work. Its covalent triple bond makes it chemically inert, safe to handle and use in proximity to reactive chemicals, while its very low boiling point (−196 °C) makes it an ideal coolant.
Covalent molecules have specific 3D shapes because electron pairs repel each other and arrange themselves as far apart as possible. This is VSEPR theory (Valence Shell Electron Pair Repulsion). Carbon dioxide (CO₂) is linear (180°) because the two double bonds point in opposite directions. Water (H₂O) is bent (104.5°) because two lone pairs on oxygen push the H–O–H angle below 109.5°. Methane (CH₄) is tetrahedral (109.5°), four identical bonds arrange themselves at equal angles.
Shape influences physical properties. Water is bent and has an uneven charge distribution, the oxygen end is slightly negative (δ−) and the hydrogen ends slightly positive (δ+). This makes water a polar molecule, meaning it interacts strongly with other polar molecules and ionic compounds, explaining why water dissolves salt. CO₂ has two polar bonds but a symmetric linear shape, the bond polarities cancel, making CO₂ a non-polar molecule that doesn't dissolve ionic compounds, which is why CO₂ fire extinguishers don't damage electronics.
Ammonia (NH₃) is pyramidal, three N–H bonds and one lone pair give a shape like a 3-legged stool. The lone pair makes NH₃ polar, which is why ammonia dissolves readily in water (producing ammonia solution) at a rate of about 700 litres per litre of water at 20 °C and 1 atm.
Water's bent shape, and the polarity that results, is the reason Sydney's drinking water supply can be treated using alum (aluminium sulfate), an ionic compound. The polar water molecules dissolve it completely, allowing the Al³⁺ ions to flocculate suspended particles and clarify billions of litres of drinking water per day.
Simple covalent molecules (H₂O, CO₂, CH₄, O₂) are small. The covalent bonds within each molecule are strong, but the forces between molecules (intermolecular forces) are weak. This is why simple molecules have low melting and boiling points, it only takes a small amount of energy to separate the molecules from each other, even though the bonds within each molecule remain intact. CO₂ sublimes at −78.5 °C because its molecules are so weakly attracted to each other.
In contrast, giant covalent structures (diamond, graphite, silicon dioxide/quartz) are not individual molecules, they are entire lattices of covalently bonded atoms extending in three dimensions with no definite end. Diamond has every carbon bonded to 4 others in a continuous 3D network. To melt diamond you must break actual covalent bonds, hence its melting point above 3500 °C. Silicon dioxide (quartz, SiO₂) is similar: each silicon bonds to 4 oxygens in a giant lattice, making quartz sand the most durable common mineral on Earth's surface.
Dry ice is solid CO₂, a simple covalent molecule that sublimes at −78.5 °C under atmospheric pressure. Diamond is giant covalent carbon, melts above 3500 °C. Both are covalent carbon compounds, but the structural scale (small molecule vs giant lattice) differs by a factor of roughly 10,000 in melting point.
Synthetic diamond, made by CSIRO researchers and companies like Element Six, is used in drill bits for Australian mineral exploration. The giant covalent structure of synthetic diamond gives it hardness of 10 Mohs, the maximum possible, allowing drill bits to cut through the hardest rock formations in WA and Queensland mines.
Simple covalent molecules like carbon dioxide have low melting points because the forces molecules are weak. The covalent bonds within each molecule are , but little energy is needed to separate the molecules. Giant covalent structures such as diamond and are lattices of atoms with no definite end. In diamond, every carbon atom is covalently bonded to others in a 3D network. Melting diamond requires breaking actual bonds, so its melting point is extremely high.
At the start of this lesson, you heard that diamond and graphite are both pure carbon, yet diamond is the hardest natural substance on Earth while graphite is soft enough to write with. That enormous difference in properties comes entirely from how the electrons are shared, not from the atoms themselves.
Now that you've worked through the lesson, how has your understanding of covalent bonding changed? Can you now explain why the type of bond, not just the elements involved, is what determines a substance's properties?
Q1. Explain how an ionic bond forms between sodium and chlorine. Include the terms 'electron transfer', 'cation', and 'anion'.
Q2. A student dissolves NaCl in water and finds it conducts electricity. They then test solid NaCl and find it does NOT conduct. Explain both results using your knowledge of ionic structure.
Q3. Evaluate the suitability of an ionic compound (e.g., NaCl) as a structural building material. Refer to at least THREE properties in your response.