3.1.1 Atomic Structure

To describe how atomic models changed, tie each piece of evidence to the feature it forced chemists to revise. A list of historical models on its own doesn’t show the reasoning.

Think evidence → model change. Scattering evidence, spectra and ionisation-energy data matter because a scientific model has to explain observations and is revised when better evidence demands it.

I’d learn the three particle values as a tiny table: proton (+1, relative mass 1), neutron (0, 1), electron (−1, about 1/1840). The electron mass is negligible on the scale used for mass number.

Keep charge and mass separate. Protons and electrons have equal-and-opposite relative charges, while protons and neutrons account for essentially all of an atom’s mass.

Picture an atom: put the protons and neutrons in the nucleus and the electrons outside it. Almost all the mass is concentrated in that tiny nucleus, while most of the atom is empty space.

The nucleus is extremely small compared with the atom, but contains essentially all of its mass. Chemical behaviour comes from the electrons outside it, particularly the outer-shell arrangement.

Keep A and Z separate when you read a nuclide symbol. Z is the proton number; A is protons + neutrons, so neutron number = A − Z. Neither changes when the species becomes an ion.

The ionic charge changes only the electron count. Two atoms with the same Z but different A values are isotopes of the same element because the proton number is unchanged.

Start with protons = Z and neutrons = A − Z. A neutral atom has Z electrons; for an ion, subtract electrons for a positive charge and add electrons for a negative charge.

For ²⁷Al³⁺, for example, Z = 13 so there are 13 protons, 27 − 13 = 14 neutrons and 10 electrons. The charge changes only the electron count.

Compare isotopes: keep the proton number fixed. The neutron number changes, so the mass number changes too; that is why they are still atoms of the same element.

So ³⁵Cl and ³⁷Cl are both chlorine because each has 17 protons. Their different neutron counts change their masses but, for neutral atoms, leave the same electron arrangement and therefore very similar chemistry.

Follow TOF in order: form positive ions, accelerate them to the same kinetic energy, let them drift through the flight tube, detect them, then convert flight time into m/z. At equal kinetic energy a lighter ion travels faster.

At the detector the positive ion gains an electron. That transfer produces a current; a larger current means more ions of that m/z have arrived. The instrument processes arrival time and current to construct the spectrum.

For the simple elemental spectra used here, the horizontal position gives m/z and the vertical intensity gives relative abundance. With mononuclear singly charged ions, the m/z values correspond directly to the isotope masses.

Read both axes before doing any calculation. Peak position tells you which isotope is present; relative peak height tells you how much of it is present.

Identify an element from the whole isotope pattern, not one peak in isolation. The combination of isotope masses and their relative abundances can provide a characteristic match.

Compare candidate elements by checking the m/z values first and then the abundance pattern. A plausible mass with the wrong isotope ratio is a poor match.

Mass spectrometry can be used to determine relative molecular mass. If the data identify a singly charged molecular ion, its m/z value corresponds numerically to Mᵣ.

Use an identified molecular-ion signal for Mᵣ. Check that the peak represents the intact molecule before reading its m/z value; another peak may come from a different ion.

Aᵣ is a weighted mean: multiply each isotope mass by its abundance, add those products and divide by the total abundance. Percentages, peak heights and simple ratios all work as long as the denominator matches the data.

Your answer must lie between the isotope masses and should sit closer to the more abundant isotope. That quick check catches a surprising number of calculator slips.

Build an electron configuration by filling the subshells in the required order, then adjust the electron count for an ion. AQA can take you up to Z = 36, so you need s, p and d subshells.

Use the periodic table as a quick check on the end of the neutral configuration. Count across the block: O is fourth across the Period 2 p block, so it has 2p⁴; in the d block use one shell number lower. Cr and Cu are the exceptions you need to check separately.

For ions, remove electrons from the highest principal energy level first; in transition-metal ions this means 4s electrons are removed before 3d. Always check the final electron total against the ionic charge.

With ionisation-energy data, look for the pattern before worrying about individual numbers. A large jump tells you the next electron would come from a lower shell; smaller changes can reveal subshell and electron-pair effects.

A very large jump after removal of n electrons suggests n outer-shell electrons. Across a period, smaller anomalies can be explained by moving into a higher-energy subshell or pairing electrons within a p orbital.

I’d learn this precisely: first ionisation energy is the enthalpy change when one mole of gaseous atoms each loses one electron to form one mole of gaseous 1+ ions.

The corresponding equation is X(g) → X⁺(g) + e⁻. “Gaseous”, “one mole” and formation of gaseous 1+ ions are all part of the definition, not optional decoration.

Write the species in the gas phase and remove exactly one electron each time. For the second ionisation, X⁺(g) → X²⁺(g) + e⁻; the charge on the ion increases by one.

Start each successive ionisation equation from the ion produced by the previous ionisation. Writing the third ionisation from X(g), for example, describes a different process.

Tie ionisation-energy patterns to electronic structure. Across Period 3 increasing nuclear charge generally raises first ionisation energy, while entering a higher-energy p subshell or pairing p electrons causes the familiar deviations. Down Group 2, extra shells and shielding make the outer electron easier to remove.

Successive ionisation energies add a second clue: a very large jump appears when removal moves from valence electrons into an inner shell. For a Group 2 atom that jump comes after the second electron, directly revealing two outer-shell electrons.