Chemical bonding is all about understanding how atoms stick together...
Chemistry426 views·Updated 11 Jul 2026·2 pagesAQA GCSE Chemistry Topic 2 Study Guide
Bella@bella_2007
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Chemical Bonding and Structure Models
Ionic bonding happens when metals meet non-metals and sparks fly! Metals lose electrons to become positively charged, whilst non-metals gain electrons to become negatively charged. These opposite charges attract strongly, forming ionic compounds with high melting points.
Here's the key thing about conducting electricity: ionic compounds only conduct when they're molten or dissolved in water. When solid, the charged particles (ions) are stuck in place, but heating frees them up to move and carry current.
Covalent bonding occurs between non-metals that share electrons rather than transferring them. Think of it as atoms holding hands by sharing electron pairs. You'll show these bonds as straight lines in structural formulae - each line represents one shared pair of electrons.
Small covalent molecules have strong bonds within each molecule but weak forces between different molecules. This means they need very little energy to melt or boil, giving them low melting and boiling points. They can't conduct electricity because they have no free-moving charged particles.
Quick Test Tip: Remember - ionic compounds conduct when liquid/dissolved, covalent molecules generally don't conduct at all!
2
of 2
Giant Structures and Nanomaterials
Diamond and graphite are both pure carbon but behave completely differently due to their structure. Diamond has each carbon bonded to four others in a rigid 3D network, making it incredibly hard. Graphite has carbon atoms in layers with only three bonds each, leaving delocalised electrons free to move and conduct electricity.
Metallic bonding creates a "sea" of delocalised electrons around metal atoms. This explains why metals conduct electricity and heat so well - those free electrons can move about easily. Pure metals are quite soft because their regular layers slide past each other, but alloys mix different-sized atoms to make sliding harder.
Polymers are huge molecules with strong covalent bonds creating long chains. Their properties depend on the forces between these chains - stronger forces mean higher melting points and tougher materials.
Nanomaterials like carbon nanotubes and fullerenes are revolutionising technology. These structures, measured in nanometres (billionths of a metre), have massive surface areas compared to their volume. This makes them brilliant catalysts and useful in everything from electronics to self-cleaning windows.
Real-World Connection: Graphene from your pencil lead could power future phone batteries because its delocalised electrons make it an excellent conductor!
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Chemistry426 views·Updated 11 Jul 2026·2 pagesAQA GCSE Chemistry Topic 2 Study Guide
Bella@bella_2007
Chemical bonding is all about understanding how atoms stick together to form different materials around you. From the salt on your chips to the graphite in your pencil, everything depends on three main types of bonding that determine whether something...
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Chemical Bonding and Structure Models
Ionic bonding happens when metals meet non-metals and sparks fly! Metals lose electrons to become positively charged, whilst non-metals gain electrons to become negatively charged. These opposite charges attract strongly, forming ionic compounds with high melting points.
Here's the key thing about conducting electricity: ionic compounds only conduct when they're molten or dissolved in water. When solid, the charged particles (ions) are stuck in place, but heating frees them up to move and carry current.
Covalent bonding occurs between non-metals that share electrons rather than transferring them. Think of it as atoms holding hands by sharing electron pairs. You'll show these bonds as straight lines in structural formulae - each line represents one shared pair of electrons.
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Giant Structures and Nanomaterials
Diamond and graphite are both pure carbon but behave completely differently due to their structure. Diamond has each carbon bonded to four others in a rigid 3D network, making it incredibly hard. Graphite has carbon atoms in layers with only three bonds each, leaving delocalised electrons free to move and conduct electricity.
Metallic bonding creates a "sea" of delocalised electrons around metal atoms. This explains why metals conduct electricity and heat so well - those free electrons can move about easily. Pure metals are quite soft because their regular layers slide past each other, but alloys mix different-sized atoms to make sliding harder.
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