Advanced Higher Chemistry Unit 1 - Summary Mind Maps

Light and Atomic Spectroscopy

Ever wondered how scientists figure out what stars are made of? It's all about light and energy levels in atoms! When you heat up elements, their electrons get excited and jump to higher energy levels - like students moving up to harder classes.

When these excited electrons fall back down, they release photons of specific colours. Each element has its own unique fingerprint of colours, which is brilliant for identification. Hydrogen produces different colours than sodium, for example.

There are two main techniques you need to know: Atomic Emission Spectroscopy (AES) measures the light given off when electrons fall down energy levels, while Atomic Absorption Spectroscopy (AAS) measures which wavelengths get absorbed. The intensity tells you how much of an element is present - dead useful for analysis!

The key equation connecting energy and light is E = hf, where h is Planck's constant (6.63×10⁻³⁴ J·s). Light behaves as both waves and particles (called photons), which might sound mental but it's absolutely fundamental to understanding atomic behaviour.

💡 Quick Tip: Remember that the further electrons are from the nucleus, the less energy difference between levels - this affects which colours you'll see!

Electronic Structure and Quantum Numbers

Your electron arrangements from earlier years were just the beginning - now we're getting into the proper detail with quantum numbers and orbital shapes! Think of orbitals as the specific "rooms" where electrons live around an atom.

S-orbitals are spherical (like footballs), p-orbitals are dumbbell-shaped, and d-orbitals have more complex shapes. Each type can hold different numbers of electrons: s holds 2, p holds 6, and d holds 10. The principal quantum number $n$ tells you the energy level.

Three crucial rules govern how electrons fill orbitals: Aufbau principle (lowest energy first), Pauli exclusion principle (maximum 2 electrons per orbital with opposite spins), and Hund's rule (fill singly before pairing up). These explain why ionisation energies don't always increase smoothly across periods.

Half-filled and completely filled subshells are particularly stable, which explains some unexpected electron configurations. For instance, chromium prefers [Ar] 3d⁵ 4s¹ rather than [Ar] 3d⁴ 4s² because the half-filled d-shell provides extra stability.

💡 Quick Tip: Use orbital box notation to visualise electron arrangements - it makes predicting magnetic properties and reactivity much clearer!

Molecular Shapes and Transition Metal Chemistry

VSEPR theory is your best mate for predicting molecular shapes - electrons hate each other and spread out as far as possible! Count up electron pairs around the central atom: 2 pairs = linear, 3 pairs = trigonal planar, 4 pairs = tetrahedral.

Lone pairs are bullies - they repel more strongly than bonding pairs, so they squash bond angles. That's why ammonia (with one lone pair) has a bond angle less than 109.5°, making it trigonal pyramidal rather than tetrahedral.

Transition metals are the colourful characters of chemistry! They form complex ions where ligands (electron pair donors) surround a central metal ion through dative bonding. The gorgeous colours come from d-d electron transitions when light hits these complexes.

These metals can form multiple oxidation states because they lose 4s electrons first, then d electrons. When forming complexes, the d-orbitals split due to ligand repulsion, creating an energy gap that corresponds perfectly to visible light wavelengths.

💡 Quick Tip: Only transition metals with partially filled d-shells show colour - if the d-shell is empty or completely full, no d-d transitions can occur!

Chemical Equilibrium

Dynamic equilibrium isn't about nothing happening - it's like a busy roundabout where cars enter and leave at exactly the same rate! The forward and reverse reaction rates become equal, but molecules are still reacting constantly.

The equilibrium constant (K) tells you who's winning the battle between reactants and products. If K > 1, products dominate; if K < 1, reactants are in charge. Temperature is the only thing that actually changes K values.

Le Chatelier's principle is your prediction tool: disturb an equilibrium and it shifts to minimise that disturbance. Add more reactant? The equilibrium shifts right to use it up. Increase pressure? It shifts towards fewer gas molecules.

ICE calculations (Initial, Change, Equilibrium) help you work out concentrations and K values. Set up a table, use algebra to track the changes, and substitute into the K expression. For temperature effects, remember: increasing temperature favours the endothermic direction.

💡 Quick Tip: Only include gases and aqueous species in K expressions - pure solids and liquids have constant concentrations so they're ignored!

Acids, Bases and pH

Forget the old "acids taste sour" definition - Brønsted-Lowry theory is where it's at! Acids are proton (H⁺) donors, bases are proton acceptors. When they react, they form conjugate acid-base pairs.

Strong acids like HCl completely dissociate in water, so pH calculations are straightforward: pH = -log[H⁺]. Weak acids only partially dissociate, so you need the acid dissociation constant Ka and the equation pH = -log Ka - log$c/[A⁻]$.

The ionic product of water $Kw = 1×10⁻¹⁴$ is absolutely crucial. It connects [H⁺] and [OH⁻] in any aqueous solution, letting you calculate pH of bases too. For strong bases, work out [OH⁻], then use Kw to find [H⁺].

Buffer solutions resist pH changes by containing a weak acid and its conjugate base. They work by neutralising small amounts of added acid or base. The Henderson-Hasselbalch equation helps calculate buffer pH.

💡 Quick Tip: Indicators change colour at their pKa value - choose one whose pKa matches your expected equivalence point pH for sharp colour changes!

Thermodynamics and Reaction Feasibility

Gibbs free energy (ΔG) is the ultimate judge of whether reactions can happen spontaneously. The equation ΔG = ΔH - TΔS combines enthalpy and entropy effects. If ΔG is negative, the reaction is feasible - products are favoured.

Entropy measures disorder, and the universe loves chaos! Solids have low entropy, gases have high entropy. When calculating ΔS, remember that reactions producing more gas molecules usually increase entropy significantly.

Temperature can make or break reaction feasibility. For endothermic reactions with positive ΔS, there's a threshold temperature where ΔG becomes negative. Set ΔG = 0 and solve: T = ΔH/ΔS gives you this crossover point.

Standard enthalpy of formation values let you calculate ΔH using Hess's law: ΔH = Σ ΔHf(products) - Σ ΔHf(reactants). Elements in their standard states have ΔHf = 0, which simplifies calculations loads.

💡 Quick Tip: Watch your units! ΔH is usually in kJ/mol but ΔS is in J/K·mol - convert ΔS by dividing by 1000 to match units in the Gibbs equation.

Reaction Kinetics

Collision theory explains why reactions have different speeds - particles need enough kinetic energy and the right orientation to react successfully. It's like trying to unlock your phone while running for the bus!

The rate-determining step (RDS) is the slowest step in any multi-step reaction mechanism. Think of it as the bottleneck that controls the overall reaction rate. The rate equation depends only on species involved in this slow step.

Reaction orders show how concentration changes affect rate. Zero order means concentration has no effect, first order means doubling concentration doubles the rate, and second order means doubling concentration quadruples the rate.

The rate constant $k$ depends on temperature and tells you how fast a reaction proceeds at given concentrations. Its units change with overall reaction order: mol⁻¹s⁻¹ for second order, s⁻¹ for first order, mol·s⁻¹ for zero order.

Rate-concentration graphs have the steepest gradient at t = 0 (the initial rate) because this is when reactant concentrations are highest. As reactions progress, rates slow down because there's less reactant left to collide.

💡 Quick Tip: You can't predict reaction orders from balanced equations - they must be determined experimentally by measuring how rate changes with concentration!