Chemical bonding is all about how atoms stick together to... Show more
Chemistry563 views·Updated May 20, 2026·6 pagesDetailed Chemistry Mind Maps for Bonding and Structure
1 / 6
1
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Metallic Bonding
Ever wondered why metals are so good at conducting electricity? It's all down to metallic bonding - one of chemistry's most fascinating concepts.
Metallic bonds form because of electrostatic attraction between positively charged metal ions and a "sea" of delocalised electrons that move freely around them. Think of it like metal atoms swimming in a pool of shared electrons that don't belong to any particular atom.
The strength of metallic bonds depends on the charge of the metal ion - the more positive the charge, the stronger the attraction and the tougher the bond. Pure metals are actually quite soft because their atoms are all the same size, so layers can slide over each other easily.
Alloys are mixtures of two or more metals that solve this problem. They're much harder than pure metals because the different-sized atoms disrupt the neat layers, preventing them from sliding past each other.
Quick Tip: Remember that delocalised electrons are the key to understanding most metallic properties!
2
of 6
Ionic Bonding
Ionic bonding happens when metals meet non-metals, and it's basically atomic give-and-take at its finest.
Here's how it works: metals transfer electrons to non-metals, creating charged particles called ions. Take sodium fluoride - sodium gives up its outer electron to fluorine, making Na⁺ and F⁻ ions. Both atoms become stable, and everyone's happy!
The magic happens through electrostatic attraction between these oppositely charged ions. It's like tiny magnets - positive attracts negative, forming strong ionic compounds.
Top tip for exams: Ionic bonding always involves at least one metal and one non-metal. Use the charge cheat sheet to work out ion charges - Group 1 metals form 1⁺ ions, Group 2 form 2⁺ ions, and so on.
Memory Hook: Think "ionic = transfer" - electrons move from one atom to another!
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Covalent Bonding
Covalent bonding is all about sharing, and it only happens between non-metal atoms who've decided to pool their electrons together.
Diatomic molecules like Cl₂ are perfect examples - two chlorine atoms share electrons to become stable. You'll see this in loads of common molecules like water (H₂O), carbon dioxide (CO₂), and ammonia (NH₃).
There are several ways to represent covalent bonds. Dot and cross diagrams show the electronic structure clearly, whilst displayed formulas use simple lines. Ball and stick models are brilliant because they show the actual 3D shape of molecules.
The key thing to remember is that electrons are shared, not transferred like in ionic bonding. This sharing creates strong bonds within molecules, but the forces between separate molecules are much weaker.
Exam Success: Practice drawing dot and cross diagrams - they're exam favourites and really help you understand what's happening!
4
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Melting and Boiling Points
Understanding melting and boiling points is crucial for predicting how different materials behave when heated.
Metallic substances have very high melting and boiling points because those delocalised electrons create incredibly strong bonds. You need loads of energy to break apart that sea of electrons and separate the metal ions.
Ionic compounds also have high melting and boiling points. The strong electrostatic attraction between positive and negative ions in the giant ionic lattice takes serious energy to overcome.
Covalent substances are the odd ones out. Whilst the bonds within molecules are strong, the intermolecular forces between separate molecules are weak. This means simple molecular substances have low melting and boiling points - you don't need much energy to separate the molecules from each other.
Key Insight: It's not about breaking the bonds within molecules, but about separating whole molecules from each other!
5
of 6
Malleability and Solubility
Malleability explains why you can bend metals without snapping them, but this property changes dramatically in alloys.
In pure metals, atoms are all the same size, so layers can slide smoothly over each other when force is applied. Alloys like steel, brass, and bronze are much harder because different-sized atoms disrupt this sliding action, making them more brittle.
Solubility is all about polar molecules, especially water. Water is polar because it has both partial positive (δ⁺) and partial negative (δ⁻) charges on different parts of the molecule.
When giant ionic substances meet water, the polar water molecules attract and separate the ions, causing the solid to dissolve. At GCSE level, remember that only giant ionic substances are soluble in water - compounds like CaCl₂, MgO, and NaOH all dissolve because they're ionic.
Water Wisdom: If it's ionic, it's likely soluble - water's polar nature is the key!
6
of 6
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Chemistry563 views·Updated May 20, 2026·6 pagesDetailed Chemistry Mind Maps for Bonding and Structure
Chemical bonding is all about how atoms stick together to form the substances around you - from the metal in your phone to the salt on your chips. There are three main types of bonding that explain why materials behave... Show more
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Metallic Bonding
Ever wondered why metals are so good at conducting electricity? It's all down to metallic bonding - one of chemistry's most fascinating concepts.
Metallic bonds form because of electrostatic attraction between positively charged metal ions and a "sea" of delocalised electrons that move freely around them. Think of it like metal atoms swimming in a pool of shared electrons that don't belong to any particular atom.
The strength of metallic bonds depends on the charge of the metal ion - the more positive the charge, the stronger the attraction and the tougher the bond. Pure metals are actually quite soft because their atoms are all the same size, so layers can slide over each other easily.
Alloys are mixtures of two or more metals that solve this problem. They're much harder than pure metals because the different-sized atoms disrupt the neat layers, preventing them from sliding past each other.
Quick Tip: Remember that delocalised electrons are the key to understanding most metallic properties!
2
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Ionic Bonding
Ionic bonding happens when metals meet non-metals, and it's basically atomic give-and-take at its finest.
Here's how it works: metals transfer electrons to non-metals, creating charged particles called ions. Take sodium fluoride - sodium gives up its outer electron to fluorine, making Na⁺ and F⁻ ions. Both atoms become stable, and everyone's happy!
The magic happens through electrostatic attraction between these oppositely charged ions. It's like tiny magnets - positive attracts negative, forming strong ionic compounds.
Top tip for exams: Ionic bonding always involves at least one metal and one non-metal. Use the charge cheat sheet to work out ion charges - Group 1 metals form 1⁺ ions, Group 2 form 2⁺ ions, and so on.
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3
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Covalent Bonding
Covalent bonding is all about sharing, and it only happens between non-metal atoms who've decided to pool their electrons together.
Diatomic molecules like Cl₂ are perfect examples - two chlorine atoms share electrons to become stable. You'll see this in loads of common molecules like water (H₂O), carbon dioxide (CO₂), and ammonia (NH₃).
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The key thing to remember is that electrons are shared, not transferred like in ionic bonding. This sharing creates strong bonds within molecules, but the forces between separate molecules are much weaker.
Exam Success: Practice drawing dot and cross diagrams - they're exam favourites and really help you understand what's happening!
4
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Melting and Boiling Points
Understanding melting and boiling points is crucial for predicting how different materials behave when heated.
Metallic substances have very high melting and boiling points because those delocalised electrons create incredibly strong bonds. You need loads of energy to break apart that sea of electrons and separate the metal ions.
Ionic compounds also have high melting and boiling points. The strong electrostatic attraction between positive and negative ions in the giant ionic lattice takes serious energy to overcome.
Covalent substances are the odd ones out. Whilst the bonds within molecules are strong, the intermolecular forces between separate molecules are weak. This means simple molecular substances have low melting and boiling points - you don't need much energy to separate the molecules from each other.
Key Insight: It's not about breaking the bonds within molecules, but about separating whole molecules from each other!
5
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Malleability and Solubility
Malleability explains why you can bend metals without snapping them, but this property changes dramatically in alloys.
In pure metals, atoms are all the same size, so layers can slide smoothly over each other when force is applied. Alloys like steel, brass, and bronze are much harder because different-sized atoms disrupt this sliding action, making them more brittle.
Solubility is all about polar molecules, especially water. Water is polar because it has both partial positive (δ⁺) and partial negative (δ⁻) charges on different parts of the molecule.
When giant ionic substances meet water, the polar water molecules attract and separate the ions, causing the solid to dissolve. At GCSE level, remember that only giant ionic substances are soluble in water - compounds like CaCl₂, MgO, and NaOH all dissolve because they're ionic.
Water Wisdom: If it's ionic, it's likely soluble - water's polar nature is the key!
6
of 6Sign up to see the content. It's free!
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What is the Knowunity AI companion?
Our AI Companion is a student-focused AI tool that offers more than just answers. Built on millions of Knowunity resources, it provides relevant information, personalised study plans, quizzes, and content directly in the chat, adapting to your individual learning journey.
Where can I download the Knowunity app?
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