Fundamental Particles and Historical Models

Scientists spent centuries piecing together what atoms actually look like inside. John Dalton thought atoms were like tiny, indivisible balls, but JJ Thomson proved him wrong by discovering electrons and proposing the "plum pudding model." Then Ernest Rutherford's famous gold foil experiment revealed that atoms have a dense nucleus, leading Niels Bohr to develop the idea of fixed electron orbitals.

Inside every atom, you'll find three key subatomic particles: protons $charge +1, mass 1$, neutrons (charge 0, mass 1), and electrons $charge -1, mass 1/1840$. The mass number (A) tells you the total protons and neutrons, whilst the atomic number (Z) gives you just the protons.

Isotopes are atoms of the same element with different numbers of neutrons - same protons, different mass. This is where mass spectrometry becomes brilliant for identifying molecules. The process involves ionisation (adding charge), acceleration (giving particles equal kinetic energy), and detection (measuring abundance).

Quick Tip: Remember that lighter particles travel faster in mass spectrometry because they all have the same kinetic energy!

Mass Spectrometry Analysis

Mass spectrometry works by accelerating ionised molecules through a magnetic field - lighter particles zip through faster because they all share the same kinetic energy. The time of flight depends on the particle's mass, making this technique perfect for identifying unknown compounds.

When you analyse a mass spectrum, the x-axis shows mass/charge ratio whilst the y-axis displays percentage abundance. The molecular ion peak (the highest mass peak) represents your original molecule and is crucial for identification.

You'll often spot smaller peaks around the main peak - these come from isotopes of the same molecule. Meanwhile, significantly lighter peaks result from fragmentation when molecules break apart inside the spectrometer.

Relative atomic mass calculations use the weighted average formula: Σ(isotope abundance × isotope mass) ÷ Σ(isotope abundance). This accounts for all isotopes naturally present in a sample.

Exam Tip: The molecular ion peak is always the furthest right peak - this gives you the molecule's actual mass!

Electron Configuration Basics

Electrons don't just randomly orbit the nucleus - they're organised in electron shells (given the symbol 'n'). The closer to the nucleus (lower n value), the lower the energy. Each shell contains sub-shells, and amazingly, a shell with value n contains exactly n sub-shells.

Sub-shells are made up of orbitals, with each orbital holding a maximum of two electrons. There are four types: s (1 orbital), p (3 orbitals), d (5 orbitals), and f (7 orbitals). These correspond to different blocks on the periodic table.

The s orbitals are spherical, whilst p orbitals have a dumbbell shape. Understanding these shapes helps explain chemical bonding later on. Each type of sub-shell can hold different numbers of electrons: s holds 2, p holds 6, d holds 10, and f holds 14.

When writing electron configurations, you'll use notation like 1s²2s²2p⁶ for neon. The superscript numbers tell you how many electrons are in each sub-shell.

Memory Trick: The periodic table is your best friend for electron configurations - just follow the blocks from left to right!

Orbital Filling Rules and Electron Spins

Getting electron configurations right means following three crucial rules. First, fill the lowest energy orbitals before moving to higher ones. Second, electrons prefer to occupy separate orbitals with the same spin before pairing up. Third, no orbital can hold more than two electrons.

Electron spins are represented by arrows pointing up or down - electrons in the same orbital must have opposite spins for stability. This is why we draw orbital diagrams with arrows showing electron arrangements.

There's a tricky bit with transition metals: the 4s sub-shell actually has lower energy than 3d, so electrons fill 4s first. However, when writing configurations for transition metal ions, electrons are removed from 4s before 3d because 3d becomes lower in energy once occupied.

Some atoms prefer unusual configurations for extra stability. Chromium, for example, adopts $Ar$3d⁵4s¹ rather than $Ar$3d⁴4s² because having all d orbitals half-filled is more stable.

Pro Tip: Always remember that 4s fills before 3d, but empties before 3d when forming ions!

Ionisation Energy Patterns

First ionisation energy is the minimum energy needed to remove one mole of electrons from one mole of gaseous atoms. It's measured in kJ/mol and tells us loads about atomic structure. The equation looks like: Na(g) → Na⁺(g) + e⁻.

Moving across a period, ionisation energy increases because atomic radius decreases and nuclear charge increases - electrons are held more tightly. Down a group, it decreases due to increased atomic radius and electron shielding reducing the nuclear pull on outer electrons.

There are some fascinating exceptions in Period 3. Aluminium has lower ionisation energy than magnesium because Al's outer electron is in a 3p orbital, which is higher energy and further from the nucleus than Mg's 3s electron.

Similarly, sulfur has lower ionisation energy than phosphorus. In sulfur's 3p⁴ configuration, two electrons must pair up in the same orbital, and their mutual repulsion makes one easier to remove compared to phosphorus's unpaired 3p electrons.

Key Insight: Exceptions in ionisation energy trends provide excellent evidence for electron sub-shells and orbital theory!

Ionisation Energy Evidence for Electron Shells

Looking at successive ionisation energies for any element reveals dramatic jumps that provide brilliant evidence for electron shells. When you plot ionisation energy against the number of electrons removed, you'll see gradual increases followed by massive leaps.

These large increases occur when you start removing electrons from a shell much closer to the nucleus. For example, removing sodium's first electron (3s¹) is relatively easy, but the second electron comes from the 2p⁶ shell - much closer to the nucleus and requiring vastly more energy.

The pattern perfectly matches our orbital theory: electrons in the same shell have similar ionisation energies, whilst electrons in inner shells require dramatically more energy to remove. This experimental evidence convinced scientists that electrons occupy distinct energy levels.

You can use these patterns to identify which group an unknown element belongs to by counting how many electrons can be removed before hitting the first major energy jump.

Exam Gold: Successive ionisation energy graphs are perfect evidence that electron shells actually exist - the jumps prove it!