Revise The nuclear atom for IB Chemistry (S1.2)

Atoms are built from three particles: protons, neutrons and electrons. Protons and neutrons, together called nucleons, sit in the small, dense, positively charged nucleus. Electrons occupy the much larger region of space around the nucleus, arranged in energy levels (shells).

To describe a specific atom precisely we use nuclide notation, which states the atomic number and mass number and so fixes the count of every particle.

Nuclide notation: $^{A}_{Z}X$

Here $X$ is the element symbol, $A$ is the mass number (protons + neutrons), and $Z$ is the atomic number (protons). Because the atomic number defines the element, $Z$ is often left off - the symbol alone tells you it. From these two numbers: protons = Z, neutrons = A − Z, and electrons = Z in a neutral atom (adjusted for any ionic charge).

Rutherford's experiment fired positive alpha particles at thin gold foil. Most passed straight through, showing the atom is mostly empty space, but a few bounced sharply back - evidence of a tiny, dense, positively charged nucleus. This is the nuclear model: protons and neutrons packed into a central nucleus, with electrons occupying the surrounding space in energy levels. It explains why the atom is electrically neutral overall and why nearly all its mass is concentrated in an almost vanishingly small volume.

All atoms of an element share the same number of protons, but they can hold different numbers of neutrons. These variants are isotopes. For example, carbon-12 ($^{12}C$) has 6 protons and 6 neutrons, while carbon-14 ($^{14}C$) has 6 protons and 8 neutrons. Isotopes react identically because they have the same electron arrangement; only their masses differ. Since most elements occur as a mixture of isotopes, we describe an element's mass with a weighted average - the relative atomic mass ($A_r$) - rather than the mass of any single atom.

$A_r = \sum (\text{isotopic mass} \times \text{fractional abundance})$

A mass spectrometer is the instrument that supplies the two numbers the $A_r$ calculation needs. In outline, a sample is vaporised and ionised, the resulting ions are accelerated and then separated according to their mass-to-charge ratio (m/z), and a detector records how many ions of each mass arrive. For a pure element, every isotope produces its own peak.

So mass spectrometry and the $A_r$ calculation are two halves of one idea: the instrument measures which isotopes are present and in what proportion, and the weighted average turns that raw abundance data into the single relative atomic mass we use in chemistry.

The nuclear atom is not just an opening topic - it underpins electron configuration, the periodic table, bonding and every mole calculation to come. Relative atomic mass, built here from isotopic abundance, is the quantity you will use to convert between mass and amount of substance throughout the course. This content is common to SL and HL at this level.