Chemistry Paper 1
AQA GCSE Chemistry Paper 1 — Higher Notes + Practicals
AQA Chemistry Paper 1 covers Topics 1–5:
Atomic structure and the periodic table
Bonding, structure and properties of matter
Quantitative chemistry
Chemical changes
Energy changes
1. Atomic Structure and the Periodic Table
Atoms
An atom contains:
Protons — charge +1, relative mass 1
Neutrons — charge 0, relative mass 1
Electrons — charge −1, relative mass almost 0
Key equations
Mass number = protons + neutrons
Atomic number = number of protons
For a neutral atom:
number of protons = number of electrons
Isotopes
Isotopes are atoms of the same element with:
The same number of protons
Different numbers of neutrons
They have similar chemical properties because they have the same number of electrons.
Electronic structure
Electrons occupy shells.
For the first 20 elements, shells fill:
2, 8, 8, 2
Examples:
Na = 2,8,1
Mg = 2,8,2
Cl = 2,8,7
Ca = 2,8,8,2
The group number tells you the number of electrons in the outer shell for Groups 1–7.
The Periodic Table
Elements are arranged by increasing atomic number.
Groups
Elements in the same group have similar chemical properties because they have the same number of outer-shell electrons.
Group 1 — Alkali metals
1 outer electron
Form +1 ions
Reactivity increases down the group
Low melting points
React with water to produce an alkaline solution and hydrogen
General reaction:
Metal + water → metal hydroxide + hydrogen
Group 7 — Halogens
7 outer electrons
Form −1 ions
Exist as diatomic molecules
React with metals to form ionic compounds
Reactivity decreases down the group
Displacement: A more reactive halogen displaces a less reactive halogen from its compound.
Group 0 — Noble gases
Full outer electron shell
Very unreactive
Exist as single atoms
Boiling points increase down the group
2. Bonding, Structure and Properties
Ionic bonding
Occurs between a metal and non-metal.
Electrons are transferred from the metal to the non-metal.
This forms oppositely charged ions.
Example:
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
NaCl forms because Na⁺ and Cl⁻ are attracted to Cl⁻.
Ionic structures
Ionic compounds form a giant ionic lattice.
They have:
High melting and boiling points
Strong electrostatic attractions
They conduct electricity when:
Molten
Dissolved in water
They do not conduct when solid because the ions cannot move.
Covalent Bonding
Occurs between non-metal atoms.
Atoms share pairs of electrons.
Simple molecular substances
Examples:
H₂
O₂
CO₂
CH₄
H₂O
They have:
Strong covalent bonds within molecules
Weak intermolecular forces between molecules
Low melting and boiling points
Usually do not conduct electricity
Important: When a molecular substance melts, you overcome intermolecular forces, not covalent bonds.
Giant Covalent Structures
Diamond
Each carbon bonded to 4 others
Very hard
Very high melting point
Does not conduct electricity
Graphite
Each carbon bonded to 3 others
Layers can slide over each other
Conducts electricity because it has delocalised electrons
Used as a lubricant and in electrodes
Graphene
One layer of graphite
Very strong
Good electrical conductor
Used in electronics and composites
Metallic Bonding
Metals consist of positive ions surrounded by delocalised electrons.
This explains why metals:
Conduct electricity
Conduct heat
Have high melting points
Are malleable
Are ductile
States of Matter
Solid
Fixed shape
Fixed volume
Particles vibrate around fixed positions
Liquid
Fixed volume
No fixed shape
Particles can move past each other
Gas
No fixed shape
No fixed volume
Particles move rapidly in random directions
Nanoparticles
Particles approximately 1–100 nm in size.
They have a very high surface area : volume ratio.
Uses include:
Sunscreens
Catalysts
Medicine
Electronics
Potential risks need to be considered because nanoparticles can behave differently from larger particles.
3. Quantitative Chemistry
Conservation of mass
In a chemical reaction:
Total mass of reactants = total mass of products
If mass appears to change, it may be because a gas has entered or escaped.
Relative Formula Mass
Mr = sum of the relative atomic masses of all atoms in the formula
Example:
H₂O:
Mr = 2(1) + 16 = 18
Moles
Key equation
Moles = mass ÷ Mr
Rearrange:
Mass = moles × Mr
Concentration
For solutions:
Concentration (g/dm³) = mass ÷ volume
For mol/dm³:
Concentration = moles ÷ volume
Remember:
1000 cm³ = 1 dm³
So:
250 cm³ = 0.250 dm³
Percentage Yield
Percentage yield = actual yield ÷ theoretical yield × 100
A yield below 100% can be caused by:
Reversible reactions
Reactants not fully reacting
Product lost during separation
Side reactions
Atom Economy
Atom economy = Mr of desired product ÷ total Mr of reactants × 100
Higher atom economy means:
Less waste
Better use of raw materials
Empirical Formula
The empirical formula shows the simplest whole-number ratio of atoms.
Method:
Find the mass of each element.
Divide each mass by its Ar.
Divide all answers by the smallest answer.
Multiply to get whole numbers if necessary.
Write the formula.
Balancing Equations
You can change the big numbers in front of formulas.
You must never change the small numbers inside formulas.
Example:
2H₂ + O₂ → 2H₂O
Reacting Masses
Typical method:
Write the balanced equation.
Calculate Mr.
Convert the known mass into moles.
Use the mole ratio from the equation.
Convert moles into the required mass.
4. Chemical Changes
Reactivity Series
Potassium
Sodium
Lithium
Calcium
Magnesium
Aluminium
Carbon
Zinc
Iron
Hydrogen
Copper
Silver
Gold
A more reactive metal can displace a less reactive metal from its compound.
Oxidation and Reduction
Remember:
OIL RIG
Oxidation Is Loss of electrons
Reduction Is Gain of electrons
Example:
Cu²⁺ + 2e⁻ → Cu
Copper ions have been reduced.
Acids
Acids produce H⁺ ions in aqueous solution.
Common acids:
Hydrochloric acid → HCl
Sulfuric acid → H₂SO₄
Nitric acid → HNO₃
pH Scale
pH < 7 = acidic
pH = 7 = neutral
pH > 7 = alkaline
A decrease of 1 pH unit means the hydrogen ion concentration increases by a factor of 10.
Acids + Metals
Acid + metal → salt + hydrogen
Example:
Mg + 2HCl → MgCl₂ + H₂
Test for hydrogen:
Lighted splint → squeaky pop
Acids + Bases
Acid + base → salt + water
This is neutralisation.
Example:
HCl + NaOH → NaCl + H₂O
Acids + Metal Oxides
Acid + metal oxide → salt + water
Example:
CuO + 2HCl → CuCl₂ + H₂O
Acids + Carbonates
Acid + carbonate → salt + water + carbon dioxide
Example:
2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂
Test for CO₂:
Bubble through limewater → turns cloudy/milky
Making a Soluble Salt
Example: making copper sulfate from copper oxide.
Warm dilute sulfuric acid.
Add copper oxide a little at a time.
Stir.
Continue adding copper oxide until some remains unreacted.
Filter to remove excess copper oxide.
Transfer the filtrate to an evaporating basin.
Heat gently to evaporate some water.
Leave to cool so crystals form.
Filter the crystals.
Dry them using filter paper.
Why add excess copper oxide?
To ensure all the acid has reacted.
Making an Insoluble Salt
Use a precipitation reaction.
Mix two soluble solutions containing the required ions.
An insoluble solid forms.
Filter the mixture.
Wash the precipitate with distilled water.
Dry the solid.
Electrolysis
Electrolysis uses electricity to break down an ionic compound.
Electrodes
Cathode = negative
Positive ions move towards the cathode.
Anode = positive
Negative ions move towards the anode.
Rules
At the cathode:
Positive ions gain electrons → reduction
At the anode:
Negative ions lose electrons → oxidation
Electrolysis of Molten Ionic Compounds
Example: molten lead bromide.
At cathode:
Pb²⁺ + 2e⁻ → Pb
At anode:
2Br⁻ → Br₂ + 2e⁻
Products:
Lead
Bromine
Electrolysis of Aqueous Solutions
In aqueous solutions, water can also provide H⁺ and OH⁻ ions.
At the cathode, hydrogen is produced if the metal is more reactive than hydrogen.
At the anode, oxygen is produced if a halide ion is not present.
5. Energy Changes
Exothermic reactions
Transfer energy to the surroundings.
Temperature increases.
Examples:
Combustion
Neutralisation
Many oxidation reactions
Energy level diagram:
Products are lower in energy than reactants.
Endothermic reactions
Take in energy from the surroundings.
Temperature decreases.
Examples:
Thermal decomposition
Some reactions involving citric acid
Energy level diagram:
Products are higher in energy than reactants.
Reaction Profiles
Activation energy
The minimum energy needed for particles to react.
A catalyst provides an alternative reaction pathway with a lower activation energy.
A catalyst:
Is not used up
Does not change the overall energy change
Increases reaction rate
Required Practicals
Practical 1: Making a Soluble Salt
Exam-style method:
Add dilute acid to a beaker.
Warm the acid gently.
Add the insoluble solid in small portions.
Stir after each addition.
Continue until some solid remains.
Filter the mixture.
Collect the filtrate.
Heat the filtrate to evaporate some water.
Leave it to cool and crystallise.
Filter the crystals.
Dry the crystals with filter paper.
Why use excess solid?
To make sure all acid has reacted.
Practical 2: Electrolysis of Aqueous Solutions
Pour the ionic solution into a suitable container.
Insert two inert electrodes.
Connect the electrodes to a DC power supply.
Switch on the power supply.
Observe the products formed at each electrode.
Test the gases where appropriate.
Record observations.
Cathode: reduction
Anode: oxidation
Practical 3: Temperature Changes
Question: Investigate the temperature change when chemicals react.
Measure a known volume of the first solution.
Measure a known volume of the second solution.
Place one solution in an insulated container.
Measure and record its initial temperature.
Add the second solution.
Stir continuously.
Record the temperature at regular intervals.
Identify the maximum or minimum temperature reached.
Repeat and calculate a mean.
For an exothermic reaction:
Temperature increases.
For an endothermic reaction:
Temperature decreases.
Improvements:
Use a better-insulated container.
Use a lid to reduce heat loss.
Use a more precise thermometer.
Repeat and calculate a mean.
Practical 4: Chromatography
This is often associated with Paper 1 chemistry practical skills.
Draw a pencil line near the bottom of chromatography paper.
Place a small spot of the sample on the line.
Allow it to dry.
Place the paper in a solvent.
Make sure the solvent level is below the pencil line.
Allow the solvent to travel up the paper.
Remove the paper before the solvent reaches the top.
Mark the solvent front.
Allow the chromatogram to dry.
Compare the spots.
Rf value
Rf = distance travelled by substance ÷ distance travelled by solvent
Rf values can be used to help identify substances.
Why pencil?
Graphite does not dissolve in the solvent, whereas ink could.
High-Priority Higher Questions
Know these extremely well:
Moles
moles = mass ÷ Mr
Concentration
concentration = moles ÷ volume in dm³
Percentage yield
actual ÷ theoretical × 100
Atom economy
Mr of desired product ÷ total Mr of reactants × 100
Percentage change
change ÷ original × 100
Rf
distance travelled by substance ÷ distance travelled by solvent
Electrolysis
Cathode = reduction
Anode = oxidation
Acids
metal → salt + hydrogen
base → salt + water
carbonate → salt + water + carbon dioxide
Gas tests
Hydrogen: squeaky pop
Oxygen: relights a glowing splint
Carbon dioxide: turns limewater cloudy
Chlorine: bleaches damp litmus paper
Practical exam rule
For almost every required practical, be ready to state:
Independent variable → what you change
Dependent variable → what you measure
Control variables → what you keep the same
Repeat → identify anomalies + calculate a mean
Accuracy → improve measuring equipment/method
Reliability → repeat results
Validity → make sure the method actually tests the intended variable