Understanding intermolecular forces, electron configuration, and acid-base reactions is crucial...
Chemistry428 views·Updated 30 Aug 2026·4 pagesOCR A-Level Chemistry Guide
Caitlin @caitlin_05
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1
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Intermolecular Forces
Ever wondered why some substances melt easily whilst others need loads of heat? It's all down to intermolecular forces - the attractions between molecules that you'll need to master for your exams.
Induced dipole-dipole forces are the weakest type and occur in all molecules. Larger molecules like iodine have bigger electron clouds, creating stronger forces and higher boiling points. That's why straight-chain hydrocarbons have higher boiling points than branched ones - they can pack together more tightly.
Permanent dipole-dipole forces happen between polar molecules like HCl. These are stronger than induced forces because the molecules have permanent positive and negative ends. You can actually test for polar molecules by seeing if a liquid stream bends towards a charged rod!
Hydrogen bonding is the strongest intermolecular force and only occurs with nitrogen, oxygen, or fluorine. This explains why water has such unusual properties - hydrogen bonds make ice less dense than liquid water, which is why ice floats.
Quick Test Tip: Remember that when heating liquids, you're breaking intermolecular forces, not covalent bonds within molecules.
2
of 4
Molecular vs Ionic Properties & Electron Configuration
The type of bonding completely determines a substance's properties. Molecular substances like iodine and ammonia have low melting points because you're only breaking weak intermolecular forces. They don't conduct electricity and their water solubility depends on polarity.
Giant ionic compounds like sodium chloride behave totally differently. They're solid at room temperature with high melting points because you need massive energy to break strong electrostatic forces between ions. They conduct when liquid because ions can move freely.
Moving on to electron configuration - this might seem tricky, but it's just about filling electron "parking spaces" in order. The s subshell holds 2 electrons, p holds 6, d holds 10, and f holds 14. Remember that electrons prefer to occupy empty orbitals first before pairing up.
For ions, simply add or remove electrons from the highest energy level. Calcium (Ca) has 20 electrons, but Ca²⁺ has lost 2 electrons, giving it 18. Use noble gas notation like [Ar] 4s¹ for potassium to save time in exams.
Memory Trick: Higher shell numbers mean electrons are further from the nucleus and have higher energy levels.
3
of 4
Orbital Shapes
Orbitals aren't just abstract concepts - they're actual 3D regions where electrons hang out around the nucleus. Understanding their shapes helps explain how atoms bond and behave.
S orbitals are perfectly spherical, whilst p orbitals are dumbbell-shaped and come in three orientations (px, py, pz). When filling orbitals, electrons spin in opposite directions to minimise repulsion - think of it like they're trying to avoid each other!
The aufbau principle is your roadmap for electron configuration. Fill lower energy levels first, and remember that 4s fills before 3d (even though 3d is written first in the final configuration). For iron (Fe), you get: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s².
When atoms form ions, they lose or gain electrons to achieve stable configurations. This is fundamental to understanding chemical reactivity and bonding patterns you'll encounter throughout A-level chemistry.
Exam Success: Always check your electron numbers add up to the atomic number for neutral atoms!
4
of 4
Acids and Bases
Acids are proton donors and bases are proton acceptors - this Brønsted-Lowry definition is absolutely essential for A-level success. When acids and bases react, they're essentially swapping protons (H⁺ ions).
Strong acids like HCl, H₂SO₄, and HNO₃ completely dissociate in water, flooding the solution with H⁺ ions. Weak acids like ethanoic acid only partially dissociate, creating an equilibrium. The same principle applies to strong bases (NaOH, KOH) versus weak bases (NH₃).
Neutralisation reactions always follow the same pattern: acid + base → salt + water. The key reaction is H⁺ + OH⁻ → H₂O. You'll also need to know that metals react with acids to produce hydrogen gas, whilst metal carbonates produce carbon dioxide.
Polyprotic acids can donate multiple protons - monoprotic (like HNO₃), diprotic (like H₂SO₄), or triprotic (like H₃PO₄). This affects stoichiometry in calculations, so always check how many protons each acid can donate.
Equation Mastery: Learn the ionic equations - they show what's actually happening chemically and often earn extra marks in exams.
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Chemistry428 views·Updated 30 Aug 2026·4 pagesOCR A-Level Chemistry Guide
Caitlin @caitlin_05
Understanding intermolecular forces, electron configuration, and acid-base reactions is crucial for success in A-level chemistry. These fundamental concepts explain everything from why water forms ice to how atoms bond and react with each other.
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Intermolecular Forces
Ever wondered why some substances melt easily whilst others need loads of heat? It's all down to intermolecular forces - the attractions between molecules that you'll need to master for your exams.
Induced dipole-dipole forces are the weakest type and occur in all molecules. Larger molecules like iodine have bigger electron clouds, creating stronger forces and higher boiling points. That's why straight-chain hydrocarbons have higher boiling points than branched ones - they can pack together more tightly.
Permanent dipole-dipole forces happen between polar molecules like HCl. These are stronger than induced forces because the molecules have permanent positive and negative ends. You can actually test for polar molecules by seeing if a liquid stream bends towards a charged rod!
Hydrogen bonding is the strongest intermolecular force and only occurs with nitrogen, oxygen, or fluorine. This explains why water has such unusual properties - hydrogen bonds make ice less dense than liquid water, which is why ice floats.
Quick Test Tip: Remember that when heating liquids, you're breaking intermolecular forces, not covalent bonds within molecules.
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Molecular vs Ionic Properties & Electron Configuration
The type of bonding completely determines a substance's properties. Molecular substances like iodine and ammonia have low melting points because you're only breaking weak intermolecular forces. They don't conduct electricity and their water solubility depends on polarity.
Giant ionic compounds like sodium chloride behave totally differently. They're solid at room temperature with high melting points because you need massive energy to break strong electrostatic forces between ions. They conduct when liquid because ions can move freely.
Moving on to electron configuration - this might seem tricky, but it's just about filling electron "parking spaces" in order. The s subshell holds 2 electrons, p holds 6, d holds 10, and f holds 14. Remember that electrons prefer to occupy empty orbitals first before pairing up.
For ions, simply add or remove electrons from the highest energy level. Calcium (Ca) has 20 electrons, but Ca²⁺ has lost 2 electrons, giving it 18. Use noble gas notation like [Ar] 4s¹ for potassium to save time in exams.
Memory Trick: Higher shell numbers mean electrons are further from the nucleus and have higher energy levels.
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Orbital Shapes
Orbitals aren't just abstract concepts - they're actual 3D regions where electrons hang out around the nucleus. Understanding their shapes helps explain how atoms bond and behave.
S orbitals are perfectly spherical, whilst p orbitals are dumbbell-shaped and come in three orientations (px, py, pz). When filling orbitals, electrons spin in opposite directions to minimise repulsion - think of it like they're trying to avoid each other!
The aufbau principle is your roadmap for electron configuration. Fill lower energy levels first, and remember that 4s fills before 3d (even though 3d is written first in the final configuration). For iron (Fe), you get: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s².
When atoms form ions, they lose or gain electrons to achieve stable configurations. This is fundamental to understanding chemical reactivity and bonding patterns you'll encounter throughout A-level chemistry.
Exam Success: Always check your electron numbers add up to the atomic number for neutral atoms!
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Acids and Bases
Acids are proton donors and bases are proton acceptors - this Brønsted-Lowry definition is absolutely essential for A-level success. When acids and bases react, they're essentially swapping protons (H⁺ ions).
Strong acids like HCl, H₂SO₄, and HNO₃ completely dissociate in water, flooding the solution with H⁺ ions. Weak acids like ethanoic acid only partially dissociate, creating an equilibrium. The same principle applies to strong bases (NaOH, KOH) versus weak bases (NH₃).
Neutralisation reactions always follow the same pattern: acid + base → salt + water. The key reaction is H⁺ + OH⁻ → H₂O. You'll also need to know that metals react with acids to produce hydrogen gas, whilst metal carbonates produce carbon dioxide.
Polyprotic acids can donate multiple protons - monoprotic (like HNO₃), diprotic (like H₂SO₄), or triprotic (like H₃PO₄). This affects stoichiometry in calculations, so always check how many protons each acid can donate.
Equation Mastery: Learn the ionic equations - they show what's actually happening chemically and often earn extra marks in exams.
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Students love us — and so will you.
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