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Structure for AQA GCSE Chemistry

Structure

This page covers the following topics:

1. Ionic compounds
2. Giant covalent structures
3. Metals and alloys

Ionic compounds form giant regular crystalline ionic lattices, held together by the strong electrostatic forces of attraction between the oppositely charged ions. Ionic lattices may be presented in 2D or 3D by including formulae of the ions or just the charge of them. Sodium chloride (NaCl) is a common example an ionic lattice and is sometimes called rock salt or table salt.

Ionic bonding is stronger than intermolecular forces in other solids. Therefore, compounds that break ionic bonds when they melt or boil, require much more energy, which results in higher melting or boiling points. In contrast, materials that only break weak intermolecular interactions while melting or boiling, require less energy and have lower melting/boiling points.

If ions from an ionic compound are free to move, a compound can conduct electricity. This is true for molten ionic compounds and ionic compoundsโ€™ solutions. Solid ionic compounds do not conduct electricity as their ions cannot move.

Ionic compounds

Some covalent compounds form giant covalent lattices that are large structures with multiple atoms joined together by covalent bonds. Compounds that form giant covalent structures break strong covalent bonds when they melt or boil, which requires a lot of energy, high temperatures and results in exceptionally high melting and boiling points.

In giant covalent structures like diamond, silicon, silicon(IV) oxide (sand, quartz), carbon and silicon atoms use all of their valence electrons to form 4 covalent bonds each. Such rigid structure makes aforementioned materials very hard but they cannot conduct electricity as there are no delocalised electrons. In fact, diamond is the hardest naturally occurring substance on Earth.

In contrast, graphite, another material made of carbon atoms, can conduct electricity and is much more fragile. Graphite consists of multiple 2D layers of carbon molecules arranged in a honeycomb structure, each of which has 3 covalent bonds, leaving remaining delocalised electrons to conduct electricity. The layers have week forces between them; and thus, they can slide easily over one another, allowing graphite to be used for pencils.

A single layer of carbon atoms arranged in a honeycomb structure found in graphite is called graphene. Graphene is the best known conductor of electricity as it also has delocalised electrons. Although it is just slightly less fragile than graphite, its strength exceeds steel hundreds of times.

Giant covalent structures

Metals form large regular structures from metal ions and delocalised electrons held together by metallic bonding. Compounds that form metallic lattices break strong metallic bonds when they melt or boil, which requires a lot of energy, high temperatures and results in high melting and boiling points. Delocalised electrons within metallic lattices can transfer electric charge and heat; thus, metals are good electrical and thermal conductors.

Due to regular structure that pure metals form, layers within the structure can relatively easily slide over one another. Less force needs to be exerted on the metal to bend it, making pure metals soft and malleable. This allows metals to be drawn into a thin wire, that is called ductility.

To introduce rigidity, different metallic elements are mixed to form alloys with preferable properties. Alloys have a more complex lattice than pure metals which results in various useful properties. For example, steel that primarily consists of iron, aluminium and carbon, is more rigid and does not rust in comparison to pure iron.

Metals and alloys

1

Why does diamond have such a high melting point?

Diamond has a giant covalent lattice consisting of silicon and oxygen atoms. Compounds that form giant covalent structures, break strong covalent bonds when they melt, which requires a lot of energy, high temperatures and results in an exceptionally high melting point.

Strong covalent bonds break when diamond melts, which requires a lot of energy, high temperature and results in a high melting point.

Why does diamond have such a high melting point?

2

What is an alloy?

Different metallic elements are mixed to form alloys.

a mixture of different metals

What is an alloy?

3

Which of the following conduct(s) electricity?

โ‹… molten ice
โ‹… molten table salt
โ‹… rock salt and water solution
โ‹… solid silver iodide precipitate
โ‹… iron oxide powder

If ions from an ionic compound are free to move, a compound can conduct electricity. This is true for molten ionic compounds and ionic compoundsโ€™ solutions. Solid ionic compounds do not conduct electricity as their ions cannot move.

โ‹… molten ice: no charged particles to transfer electricity โ†’ no
โ‹… molten table salt: ions that are free to move โ†’ yes
โ‹… rock salt and water solution: ions that are free to move โ†’ yes
โ‹… solid silver iodide precipitate: ions that are not free to move โ†’ no
โ‹… iron oxide powder: ions that are not free to move โ†’ no

molten table salt, rock salt and water solution

Which of the following conduct(s) electricity? 

โ‹… molten ice 
โ‹… molten table salt 
โ‹… rock salt and water solution 
โ‹… solid silver iodide precipitate 
โ‹… iron oxide powder

4

Why does pure lithium chloride only conduct electricity above its melting point?

If ions from an ionic compound are free to move, a compound can conduct electricity. This is true for molten ionic compounds and ionic compoundsโ€™ solutions. Solid ionic compounds do not conduct electricity as their ions cannot move. When lithium chloride reaches its melting point it melts, its ions are free to move; and thus, it conducts electricity.

When lithium chloride reaches its melting point it melts, its ions are free to move; and thus, it conducts electricity.

Why does pure lithium chloride only conduct electricity above its melting point?

5

Explain why diamond cannot conduct electricity.

Diamond consists of carbon atoms arranged in a giant covalent structure. In giant covalent structures like diamond carbon atoms use all of their valence electrons to form 4 covalent bonds each. Such rigid structure cannot conduct electricity as there are no delocalised electrons.

Diamond has no free moving ions or electrons that could carry the charge.

Explain why diamond cannot conduct electricity.

End of page

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