Chemistry Cheatsheet Pack Cover

This is your ultimate A-Level AQA Chemistry revision resource - a comprehensive cheatsheet pack designed to condense the entire course into easily digestible pages. Each section focuses on the most important definitions, key terms, diagrams and concepts you'll need to succeed in your exams.

The pack has been carefully structured to match the AQA specification points, ensuring you won't miss any crucial information. From atomic structure to organic synthesis, every topic is covered with clear explanations and practical examples.

Quick Tip: Use this pack alongside your textbooks for focused revision - it's perfect for last-minute cramming or reinforcing your understanding of complex topics.

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How to Use This Pack

This cheatsheet pack transforms the entire A-Level Chemistry course into super condensed pages that focus on what really matters. You'll find the most important definitions, key terms, diagrams and essential concepts all in one place.

The content maps perfectly to AQA specification points from 3.1.1 through 3.3.16. This covers everything from atomic structure and mass spectrometry to NMR and chromatography. Each topic has been carefully selected to ensure you don't miss anything crucial for your exams.

The pack includes organic chemistry fundamentals like alkanes, haloalkanes, alkenes and alcohols. You'll also find advanced topics such as transition metals, equilibrium constants, and protein structures. The comprehensive coverage means this can be your go-to revision resource throughout your A-Level journey.

Pro Tip: This pack works brilliantly for last-minute revision or as a quick reference guide whilst doing practice papers!

Atomic Structure, Mass, Mass Spectrometry & Electronic Configuration

Understanding atomic structure is absolutely fundamental to everything else in chemistry. Atoms consist of protons, neutrons and electrons, with the nucleus containing most of the mass whilst electrons orbit in shells around it.

The atomic model has evolved significantly over time. Dalton's indivisible atoms gave way to Thomson's plum pudding model, then Rutherford discovered the nucleus, and finally Bohr proposed electron shells. Today's model recognises that protons and neutrons are made of smaller particles called quarks.

Mass spectrometry is brilliant for identifying unknown compounds and determining isotopic abundance. The time-of-flight method involves ionisation, acceleration, drift and detection. You can calculate relative atomic mass using the formula: Ar = Σ(relative isotopic mass × abundance)/100.

Electronic configuration follows specific rules - electrons fill the lowest energy levels first (Aufbau principle), with orbitals of the same energy filling singly before pairing. Remember that 4s fills before 3d because it's actually lower in energy.

Remember: Isotopes have the same number of protons but different numbers of neutrons - they react identically but have different physical properties!

Ionisation Energy & Basic Equations

Ionisation energy measures how much energy you need to remove electrons from atoms. The first ionisation energy removes one electron to form a 1+ ion, and successive ionisation energies get progressively larger as you remove more electrons.

Three key factors affect ionisation energies: atomic radius (larger atoms hold electrons less tightly), nuclear charge (more protons mean stronger attraction), and shielding (inner electrons repel outer ones). This explains why ionisation energy increases across periods but decreases down groups.

The mole is chemistry's way of counting particles - one mole contains Avogadro's number (6.022 × 10²³) of particles. Key equations include n = m/M for calculating moles from mass, and c = n/V for concentration calculations.

Empirical formulas show the simplest whole-number ratio of atoms in compounds. You calculate them by finding moles of each element, then dividing by the smallest number to get the ratio. Molecular formulas are whole-number multiples of empirical formulas.

Atom economy and percentage yield measure reaction efficiency. Atom economy considers how much of your starting materials ends up in the desired product, whilst percentage yield compares actual to theoretical yields.

Key Insight: Successive ionisation energies show clear jumps between electron shells - this provides excellent evidence for atomic structure!

Bonds

Ionic bonding happens when metals transfer electrons to non-metals, creating oppositely charged ions that attract electrostatically. The resulting lattice structures are giant 3D arrangements where each ion is surrounded by oppositely charged neighbours.

Ionic bond strength depends on two factors: charge on the ions (higher charges mean stronger bonds) and size of the ions (smaller ions get closer together, creating stronger attractions). This explains why MgO has a much higher melting point than NaCl.

Covalent bonds form when non-metals share electron pairs to achieve stable outer shells. You can have single, double or triple bonds depending on how many electron pairs are shared. Coordinate bonds occur when one atom donates both electrons in the shared pair.

Different crystal structures have completely different properties. Ionic crystals conduct electricity when molten and dissolve in polar solvents. Covalent networks like diamond are incredibly hard with high melting points, whilst molecular crystals are much softer with lower melting points.

Quick Check: Remember that ionic compounds conduct electricity when molten or dissolved because the ions become free to move!

Enthalpies, Calorimetry & Hess's Law

Enthalpy is the heat energy stored in chemical systems, and enthalpy changes tell you whether reactions release or absorb energy. Exothermic reactions have negative ΔH values and warm up the surroundings, whilst endothermic reactions have positive ΔH values and cool things down.

Calorimetry lets you measure enthalpy changes experimentally using the equation q = mcΔT. Coffee cup calorimeters work well for neutralisation reactions, whilst spirit burner calorimeters measure combustion enthalpies. Remember that experimental values often differ from data book values due to heat losses and incomplete reactions.

Hess's Law states that enthalpy change is independent of the route taken - this is incredibly useful for calculating enthalpies you can't measure directly. You can use either combustion data or formation data to construct Hess cycles and find unknown enthalpy changes.

Bond enthalpies provide another way to estimate enthalpy changes. Bond breaking is always endothermic whilst bond making is exothermic. The equation ΔH = Σ(bonds broken) - Σ(bonds made) gives approximate values, though they're less accurate than Hess cycle calculations.

Lab Tip: When doing calorimetry experiments, plot temperature vs time and extrapolate back to find the true temperature change!

Collision Theory, Maxwell-Boltzmann, Rate of Reaction & Chemical Equilibrium

Collision theory explains why reactions happen - particles must collide with sufficient energy (above the activation energy) and in the correct orientation. Most collisions don't result in reactions because these conditions aren't met.

The Maxwell-Boltzmann distribution shows how molecular energies are spread in a gas. Only molecules with energy above Ea can react, and the curve never starts from zero because no molecules have zero energy. Temperature increases shift the whole curve to higher energies.

Catalysts speed up reactions by providing alternative pathways with lower activation energies. Homogeneous catalysts are in the same phase as reactants, whilst heterogeneous catalysts work through adsorption and desorption on their surfaces.

Chemical equilibrium occurs in closed systems when forward and backward reaction rates are equal, keeping concentrations constant. Le Chatelier's principle predicts how equilibrium responds to changes - the system always opposes the change you make.

Temperature, pressure, and concentration changes all shift equilibrium position, but catalysts only affect the rate at which equilibrium is reached. Understanding these principles is crucial for optimising industrial processes.

Key Point: Increasing temperature favours the endothermic direction, whilst increasing pressure favours the side with fewer gas molecules!

The Haber Process, Equilibrium Constant, Redox Reactions & Lattice Enthalpy

The Haber process demonstrates real-world applications of equilibrium principles. N₂ + 3H₂ ⇌ 2NH₃ with ΔH = -92 kJ mol⁻¹. Theoretically, you'd want high pressure and low temperature for maximum yield, but these aren't economically viable.

Industrial compromise conditions balance yield with practicality: 400-500°C gives reasonable rates, 200 atm provides good yield without excessive costs, and iron catalysts allow lower temperatures. These decisions involve economic and safety considerations.

The equilibrium constant Kc shows where equilibrium lies. When Kc > 1, products are favoured; when Kc < 1, reactants dominate. Temperature affects Kc values, but pressure, concentration and catalysts don't change the constant itself.

Redox reactions involve electron transfer - oxidation is electron loss (increase in oxidation number) whilst reduction is electron gain (decrease in oxidation number). You can combine half-equations by balancing electrons first, then combining and cancelling.

Lattice enthalpy measures ionic bond strength but can't be measured directly. You need other enthalpy values like atomisation, ionisation, and electron affinity to calculate it using energy cycles.

Industrial Chemistry: The Haber process shows how theory meets reality - compromise conditions are essential for profitable production!

Born-Haber Cycles, Entropy & Gibbs Free Energy

Born-Haber cycles use Hess's Law to calculate lattice enthalpies indirectly. You construct energy level diagrams showing all the steps from elements to ionic compounds, including atomisation, ionisation, electron affinity, and formation enthalpies.

These cycles become more complex for Group 2 compounds because you need second ionisation energies and multiple electron affinities. Differences between theoretical and experimental lattice enthalpies suggest covalent character in supposedly ionic bonds.

Entropy measures disorder in systems - solids have low entropy, gases have high entropy. The universe's total entropy always increases in spontaneous processes. You can predict entropy changes by counting gas molecules in reactions.

Gibbs free energy combines enthalpy and entropy effects: ΔG = ΔH - TΔS. Reactions are only feasible when ΔG ≤ 0, which depends on temperature. Some reactions become feasible only above certain temperatures when the TΔS term dominates.

The key insight is that thermodynamic feasibility doesn't guarantee that reactions will actually happen - kinetic factors like activation energy and reaction rates also matter enormously.

Top Tip: Remember that feasibility (ΔG < 0) doesn't mean a reaction will happen quickly - kinetics and thermodynamics are completely different considerations!