How does the way particles are arranged in a solid explain its melting point, hardness and conductivity?
The four types of crystalline solid (ionic, simple molecular, giant covalent and metallic), their structures, and how structure and bonding explain physical properties.
An Eduqas A-Level Chemistry C1.5 answer on ionic, simple molecular, giant covalent and metallic crystalline solids and how their structure and bonding explain melting point, hardness, solubility and conductivity.
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What this topic is asking
Eduqas topic C1.5 covers the four classes of crystalline solid and the way their structure and bonding determine physical properties such as melting temperature, hardness, electrical conductivity and solubility. The skill rewarded is explaining a property by naming the particles, the forces between them and what happens to those forces when the solid is heated, dissolved or stressed.
Ionic crystals
Simple molecular crystals
In a simple molecular solid, discrete molecules sit at lattice points held by weak intermolecular forces (van der Waals, dipole-dipole, or hydrogen bonds). The covalent bonds within molecules are strong, but the forces between them are weak, so these solids have low melting points and do not conduct (no charged particles are free to move).
Giant covalent crystals
Diamond and graphite are the contrasting allotropes of carbon: diamond bonds every atom four times (hard, insulating), while graphite bonds each atom three times in layers with delocalised electrons (soft, slippery, conducting along the layers).
Metallic crystals
A metallic solid is a regular lattice of positive ions in a sea of delocalised electrons. The strong attraction between ions and electrons gives high melting points; the mobile electrons make metals good electrical and thermal conductors; and layers of ions can slide without breaking the bonding, so metals are malleable and ductile.
Explaining properties
Eduqas marks reward a clear chain: name the particles, name the force between them, and say what heating or stressing does to that force. "It has strong bonds" alone scores nothing; "many strong ionic attractions in the giant lattice must be overcome, requiring much energy" scores the mark.
Examples in context
Example 1. Silicon dioxide in glass and sand. Its giant covalent network of Si-O bonds gives a very high melting point and great hardness, which is why quartz is used as an abrasive and why glass must be heated strongly to be worked.
Example 2. Alloys harden metals. Adding atoms of a different size (as in steel or bronze) disrupts the regular layers of the metal lattice, making it harder for layers to slide, which is why alloys are usually harder and stronger than the pure metal.
Try this
Q1. Explain why simple molecular substances such as iodine have low melting points. [2 marks]
- Cue. The molecules are held by weak van der Waals (intermolecular) forces, which need little energy to overcome; the strong covalent bonds within the molecules are not broken on melting.
Q2. State and explain one property that makes copper suitable for electrical wiring. [2 marks]
- Cue. It conducts electricity because its delocalised electrons are free to move; it is also ductile because layers of ions can slide, so it can be drawn into wires.
Exam-style practice questions
Practice questions written in the style of WJEC Eduqas exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
Eduqas 20194 marksDiamond and graphite are both forms of carbon. (a) Describe the structure and bonding in graphite. (b) Explain why graphite conducts electricity but diamond does not.Show worked answer →
(a) Graphite is a giant covalent structure of carbon atoms in layers of hexagonal rings; each carbon forms three covalent bonds, leaving one delocalised electron, and the layers are held together by weak van der Waals forces (2).
(b) Graphite conducts because each carbon contributes one delocalised electron that is free to move along the layers and carry charge (1). In diamond every carbon forms four covalent bonds with no spare electrons, so there are no free charge carriers and it does not conduct (1).
Eduqas 20203 marksExplain why magnesium oxide has a much higher melting temperature than sodium chloride, even though both are ionic solids.Show worked answer →
Magnesium oxide contains and ions, both doubly charged, while sodium chloride contains singly charged and (1).
The greater ionic charges (and smaller ionic radii) in MgO give much stronger electrostatic attraction in the lattice, so the lattice enthalpy is larger (1). More energy is therefore needed to overcome the attractions and melt the solid, giving MgO a much higher melting temperature (1).
Related dot points
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Sources & how we know this
- WJEC Eduqas GCE A Level Chemistry specification (from 2015) — WJEC Eduqas (2015)