Chemistry might seem overwhelming at first, but it's actually about...
Chemistry202 views·Updated 11 Sept 2026·9 pagesGrade 9 AQA Trilogy Higher Science: Chemistry Paper 1 Mind Maps
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1
of 9
Atoms and Chemical Reactions
Everything you see is made of atoms - the smallest parts of elements that can't be broken down further. When different elements join together, they form compounds with completely new properties.
Chemical formulae tell you exactly what's in a compound and how many atoms of each element are present. In chemical equations, reactants go on the left side and products on the right. Here's the key rule: the total mass never changes during a reaction - what goes in must equal what comes out.
Mixtures are different from compounds because the substances aren't chemically bonded. You can separate mixtures using physical processes like filtration (for separating solids), crystallisation (for getting pure salts), and chromatography (for separating coloured substances).
Quick Tip: Remember that in reactions, atoms are just rearranged - they're never created or destroyed, which is why mass stays constant.
Scientists' understanding of atoms has evolved dramatically. Thomson discovered electrons and proposed the "plum pudding" model, but Rutherford's gold foil experiment showed that atoms have a tiny, dense nucleus. Today we know atoms contain protons (positive), neutrons (neutral), and electrons (negative) arranged in specific ways.
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of 9
The Periodic Table and Group Properties
The periodic table wasn't built overnight. Newlands tried first in 1864, but Mendeleev cracked it in 1869 by leaving gaps for undiscovered elements and arranging them by properties rather than just atomic weight.
Group 0 (noble gases) are the ultimate loners - they've got full outer electron shells, making them incredibly stable and unreactive. Group 1 (alkali metals) are the complete opposite. With just one outer electron, they're desperate to react and become more reactive as you go down the group.
These alkali metals react violently with water, producing metal hydroxides and hydrogen gas. That's why lithium, sodium, and potassium are stored under oil - they'd react explosively with air or water otherwise.
Remember: Reactivity increases down Group 1 because the outer electron gets further from the nucleus, making it easier to lose.
Group 7 (halogens) are non-metals that love gaining electrons. They become less reactive going down the group and can displace less reactive halogens from solutions. Fluorine is the most reactive, while iodine is the least.
3
of 9
Molecular Structures and Carbon Allotropes
Small molecules like water and chlorine are held together by strong covalent bonds within the molecule, but only weak forces between different molecules. This is why they often have low melting and boiling points.
Carbon is incredibly versatile, forming different structures called allotropes. Diamond has each carbon atom bonded to four others in a giant 3D structure, making it incredibly hard. Graphite has carbon atoms arranged in layers that can slide over each other, making it soft and slippery.
Graphene is essentially a single layer of graphite - it's amazingly strong yet nearly transparent. Fullerenes are hollow carbon structures that can be spheres or tubes (nanotubes), useful for drug delivery and electronics.
Key Point: The same element (carbon) can have completely different properties depending on how its atoms are arranged.
The key to understanding why these materials behave so differently lies in their bonding. Graphite conducts electricity because of delocalised electrons that can move freely, while diamond doesn't conduct because all electrons are locked in covalent bonds.
4
of 9
States of Matter and Bonding
The particle model explains how solids, liquids, and gases behave differently. In solids, particles vibrate in fixed positions. In liquids, they move around each other. In gases, they zoom about randomly with loads of space between them.
Ionic bonding happens when metals give electrons to non-metals, creating charged ions. The resulting compounds have high melting points because of strong electrostatic forces between oppositely charged ions. They only conduct electricity when molten or dissolved because the ions need to be free to move.
Covalent bonding involves sharing electrons between non-metal atoms. You can have single bonds (sharing one pair) or double bonds (sharing two pairs). The number of bonds depends on how many electrons each atom needs for a full outer shell.
Essential: Ionic compounds conduct when liquid but not solid; covalent compounds generally don't conduct at all.
Giant covalent structures like diamond have billions of atoms all joined by strong covalent bonds. This gives them extremely high melting points because you need massive amounts of energy to break all those bonds.
5
of 9
Chemical Calculations and Moles
Avogadro's constant (6.02 × 10²³) tells you how many particles are in one mole of any substance. One mole of a substance has a mass equal to its relative formula mass (Mr) in grams.
The key formula you'll use constantly is: moles = mass ÷ Mr. In balanced equations, you can work out what mass of products you'll get from any given mass of reactants by converting everything to moles first.
Limiting reactants are the substances that run out first in a reaction. To find them, calculate how much you need of each reactant and compare it to what you've actually got. The one you haven't got enough of limits how much product you can make.
Exam Tip: Always balance your equation first, then convert masses to moles, then use the equation ratios.
For gases, one mole occupies 24 dm³ at room temperature and pressure. Concentration is measured in mol/dm³ and equals moles of solute divided by volume of solution. These calculations link together to help you work out exactly what happens in reactions.
6
of 9
Metal Reactivity and Extraction
The reactivity series ranks metals from most reactive (potassium) to least reactive (gold). More reactive metals react more vigorously with water and acids, often producing hydrogen gas and metal compounds.
Displacement reactions happen when a more reactive metal kicks out a less reactive metal from its compound. For example, iron displaces copper from copper sulfate solution because iron is more reactive than copper.
Metal extraction depends on reactivity. Unreactive metals like gold are found pure in the ground. More reactive metals are found as ores (metal compounds) and need to be extracted using methods like heating with carbon or electrolysis.
Key Rule: More reactive metals displace less reactive metals from their compounds.
Oxidation means gaining oxygen or losing electrons. Reduction means losing oxygen or gaining electrons. In displacement reactions, the more reactive metal gets oxidised while the metal ion in the compound gets reduced. You can write these as half equations showing electron transfer.
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of 9
Acids, Bases and Salt Preparation
Acids release H⁺ ions when dissolved in water. The three main ones are hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃). They react with metals, metal oxides, hydroxides, and carbonates to form salts.
Bases neutralise acids to form water. Metal hydroxides that dissolve in water are called alkalis and release OH⁻ ions. The pH scale shows how acidic or alkaline something is - below 7 is acidic, above 7 is alkaline, and 7 is neutral.
The general neutralisation reaction is: H⁺(aq) + OH⁻(aq) → H₂O. Acids react with carbonates to produce salt, water, and carbon dioxide. With some metals, they produce salt and hydrogen gas instead.
Practical Tip: To make pure salt crystals, react excess base with acid, filter off the excess, then evaporate the water slowly.
For crystallisation practicals, you heat the acid, add excess insoluble base until no more reacts, filter to remove excess, then evaporate the solution slowly to form crystals. This gives you pure salt crystals every time.
8
of 9
Energy Changes in Reactions
Chemical reactions either give out energy (exothermic) or take in energy (endothermic). Exothermic reactions like combustion and neutralisation make the surroundings warmer. Endothermic reactions like thermal decomposition make the surroundings cooler.
You can spot the difference with a thermometer - temperature goes up for exothermic reactions and down for endothermic ones. Hand warmers and self-heating cans use exothermic reactions, while sports injury packs use endothermic reactions.
Energy is always conserved in reactions. The energy change depends on the difference between energy needed to break bonds (endothermic) and energy released when making new bonds (exothermic).
Memory Aid: Breaking bonds needs energy (endothermic), making bonds releases energy (exothermic).
To calculate energy changes using bond energies, add up all the energy needed to break reactant bonds, then subtract all the energy released making product bonds. A positive result means endothermic, negative means exothermic.
9
of 9
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Chemistry202 views·Updated 11 Sept 2026·9 pagesGrade 9 AQA Trilogy Higher Science: Chemistry Paper 1 Mind Maps
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Chemistry might seem overwhelming at first, but it's actually about understanding the building blocks of everything around you. This guide covers the fundamentals - from atoms and molecules to chemical reactions and energy changes - giving you the essential knowledge...
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Atoms and Chemical Reactions
Everything you see is made of atoms - the smallest parts of elements that can't be broken down further. When different elements join together, they form compounds with completely new properties.
Chemical formulae tell you exactly what's in a compound and how many atoms of each element are present. In chemical equations, reactants go on the left side and products on the right. Here's the key rule: the total mass never changes during a reaction - what goes in must equal what comes out.
Mixtures are different from compounds because the substances aren't chemically bonded. You can separate mixtures using physical processes like filtration (for separating solids), crystallisation (for getting pure salts), and chromatography (for separating coloured substances).
Quick Tip: Remember that in reactions, atoms are just rearranged - they're never created or destroyed, which is why mass stays constant.
Scientists' understanding of atoms has evolved dramatically. Thomson discovered electrons and proposed the "plum pudding" model, but Rutherford's gold foil experiment showed that atoms have a tiny, dense nucleus. Today we know atoms contain protons (positive), neutrons (neutral), and electrons (negative) arranged in specific ways.
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The Periodic Table and Group Properties
The periodic table wasn't built overnight. Newlands tried first in 1864, but Mendeleev cracked it in 1869 by leaving gaps for undiscovered elements and arranging them by properties rather than just atomic weight.
Group 0 (noble gases) are the ultimate loners - they've got full outer electron shells, making them incredibly stable and unreactive. Group 1 (alkali metals) are the complete opposite. With just one outer electron, they're desperate to react and become more reactive as you go down the group.
These alkali metals react violently with water, producing metal hydroxides and hydrogen gas. That's why lithium, sodium, and potassium are stored under oil - they'd react explosively with air or water otherwise.
Remember: Reactivity increases down Group 1 because the outer electron gets further from the nucleus, making it easier to lose.
Group 7 (halogens) are non-metals that love gaining electrons. They become less reactive going down the group and can displace less reactive halogens from solutions. Fluorine is the most reactive, while iodine is the least.
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Molecular Structures and Carbon Allotropes
Small molecules like water and chlorine are held together by strong covalent bonds within the molecule, but only weak forces between different molecules. This is why they often have low melting and boiling points.
Carbon is incredibly versatile, forming different structures called allotropes. Diamond has each carbon atom bonded to four others in a giant 3D structure, making it incredibly hard. Graphite has carbon atoms arranged in layers that can slide over each other, making it soft and slippery.
Graphene is essentially a single layer of graphite - it's amazingly strong yet nearly transparent. Fullerenes are hollow carbon structures that can be spheres or tubes (nanotubes), useful for drug delivery and electronics.
Key Point: The same element (carbon) can have completely different properties depending on how its atoms are arranged.
The key to understanding why these materials behave so differently lies in their bonding. Graphite conducts electricity because of delocalised electrons that can move freely, while diamond doesn't conduct because all electrons are locked in covalent bonds.
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States of Matter and Bonding
The particle model explains how solids, liquids, and gases behave differently. In solids, particles vibrate in fixed positions. In liquids, they move around each other. In gases, they zoom about randomly with loads of space between them.
Ionic bonding happens when metals give electrons to non-metals, creating charged ions. The resulting compounds have high melting points because of strong electrostatic forces between oppositely charged ions. They only conduct electricity when molten or dissolved because the ions need to be free to move.
Covalent bonding involves sharing electrons between non-metal atoms. You can have single bonds (sharing one pair) or double bonds (sharing two pairs). The number of bonds depends on how many electrons each atom needs for a full outer shell.
Essential: Ionic compounds conduct when liquid but not solid; covalent compounds generally don't conduct at all.
Giant covalent structures like diamond have billions of atoms all joined by strong covalent bonds. This gives them extremely high melting points because you need massive amounts of energy to break all those bonds.
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Chemical Calculations and Moles
Avogadro's constant (6.02 × 10²³) tells you how many particles are in one mole of any substance. One mole of a substance has a mass equal to its relative formula mass (Mr) in grams.
The key formula you'll use constantly is: moles = mass ÷ Mr. In balanced equations, you can work out what mass of products you'll get from any given mass of reactants by converting everything to moles first.
Limiting reactants are the substances that run out first in a reaction. To find them, calculate how much you need of each reactant and compare it to what you've actually got. The one you haven't got enough of limits how much product you can make.
Exam Tip: Always balance your equation first, then convert masses to moles, then use the equation ratios.
For gases, one mole occupies 24 dm³ at room temperature and pressure. Concentration is measured in mol/dm³ and equals moles of solute divided by volume of solution. These calculations link together to help you work out exactly what happens in reactions.
6
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Metal Reactivity and Extraction
The reactivity series ranks metals from most reactive (potassium) to least reactive (gold). More reactive metals react more vigorously with water and acids, often producing hydrogen gas and metal compounds.
Displacement reactions happen when a more reactive metal kicks out a less reactive metal from its compound. For example, iron displaces copper from copper sulfate solution because iron is more reactive than copper.
Metal extraction depends on reactivity. Unreactive metals like gold are found pure in the ground. More reactive metals are found as ores (metal compounds) and need to be extracted using methods like heating with carbon or electrolysis.
Key Rule: More reactive metals displace less reactive metals from their compounds.
Oxidation means gaining oxygen or losing electrons. Reduction means losing oxygen or gaining electrons. In displacement reactions, the more reactive metal gets oxidised while the metal ion in the compound gets reduced. You can write these as half equations showing electron transfer.
7
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Acids, Bases and Salt Preparation
Acids release H⁺ ions when dissolved in water. The three main ones are hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃). They react with metals, metal oxides, hydroxides, and carbonates to form salts.
Bases neutralise acids to form water. Metal hydroxides that dissolve in water are called alkalis and release OH⁻ ions. The pH scale shows how acidic or alkaline something is - below 7 is acidic, above 7 is alkaline, and 7 is neutral.
The general neutralisation reaction is: H⁺(aq) + OH⁻(aq) → H₂O. Acids react with carbonates to produce salt, water, and carbon dioxide. With some metals, they produce salt and hydrogen gas instead.
Practical Tip: To make pure salt crystals, react excess base with acid, filter off the excess, then evaporate the water slowly.
For crystallisation practicals, you heat the acid, add excess insoluble base until no more reacts, filter to remove excess, then evaporate the solution slowly to form crystals. This gives you pure salt crystals every time.
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Energy Changes in Reactions
Chemical reactions either give out energy (exothermic) or take in energy (endothermic). Exothermic reactions like combustion and neutralisation make the surroundings warmer. Endothermic reactions like thermal decomposition make the surroundings cooler.
You can spot the difference with a thermometer - temperature goes up for exothermic reactions and down for endothermic ones. Hand warmers and self-heating cans use exothermic reactions, while sports injury packs use endothermic reactions.
Energy is always conserved in reactions. The energy change depends on the difference between energy needed to break bonds (endothermic) and energy released when making new bonds (exothermic).
Memory Aid: Breaking bonds needs energy (endothermic), making bonds releases energy (exothermic).
To calculate energy changes using bond energies, add up all the energy needed to break reactant bonds, then subtract all the energy released making product bonds. A positive result means endothermic, negative means exothermic.
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We thought you’d never ask...
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.
You can download the app from Google Play Store and Apple App Store.
That's right! Enjoy free access to study content, connect with fellow students, and get instant help – all at your fingertips.
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