Atomic structure: from subatomic particles to isotopes

Almost everything else in SPM Chemistry rests on the atom. Bonding, the periodic table, formulae, reactions, all of it is really about how atoms are built and how their electrons behave. So the atomic structure part of the Matter and the Atomic Structure chapter is worth getting genuinely solid, not just passable. This explainer walks from the three subatomic particles up to isotopes and relative atomic mass, in the order the ideas build on each other.

The three subatomic particles

An atom has a tiny, dense nucleus at its centre containing protons and neutrons, with electrons moving around it in shells. What you must remember is their relative mass and charge:

• Proton, relative mass 1, charge +1, found in the nucleus.
• Neutron, relative mass 1, charge 0 (neutral), found in the nucleus.
• Electron, relative mass about 1/1840 (effectively negligible), charge −1, moving in shells around the nucleus.

Two consequences follow immediately. First, because electrons are so light, almost all the mass of an atom is in the nucleus. Second, because an atom overall is electrically neutral, the number of protons equals the number of electrons in a neutral atom.

Proton number and nucleon number

These two numbers describe any atom completely, and mixing them up is the most common early mistake.

• The proton number (Z) is the number of protons in the nucleus. It is what defines the element, every carbon atom has proton number 6, every oxygen atom has proton number 8. Change the proton number and you have a different element.
• The nucleon number (A) is the total number of protons and neutrons in the nucleus.

From these you can get everything: number of neutrons = nucleon number − proton number, and in a neutral atom number of electrons = proton number. For example, an atom with proton number 11 and nucleon number 23 (sodium) has 11 protons, 11 electrons, and 23 − 11 = 12 neutrons.

Electron arrangement

Electrons occupy shells around the nucleus, filling from the inside out. For the first 20 elements the shells hold a maximum of 2, then 8, then 8 electrons. So sodium’s 11 electrons arrange as 2, 8, 1, two shells full and one electron in the third. The valence electrons are those in the outermost shell (one, for sodium), and they control how the atom bonds and where it sits in the periodic table. This is the bridge from atomic structure straight into the periodic table: the group number matches the valence electrons, the period number matches the number of occupied shells.

Isotopes

Here is the idea students most often state loosely. Isotopes are atoms of the same element with the same proton number but different numbers of neutrons, and therefore different nucleon numbers. Because the proton number (and so the number of electrons) is unchanged, isotopes have identical chemical properties; they differ only in mass and in physical properties that depend on mass, such as density.

The standard examples:

• Chlorine-35 and chlorine-37, both have 17 protons, but 18 and 20 neutrons.
• Carbon-12, carbon-13 and carbon-14, all 6 protons, with 6, 7 and 8 neutrons.

Note the notation: the number written after the name is the nucleon number, not the number of neutrons.

Relative atomic mass, a worked example

Because a natural sample of an element is a mixture of its isotopes, we use a relative atomic mass, the average mass of its atoms compared with 1/12 the mass of a carbon-12 atom, weighted by how abundant each isotope is.

Take chlorine, which is roughly 75% chlorine-35 and 25% chlorine-37. The relative atomic mass is the weighted average:

(75 × 35 + 25 × 37) ÷ 100 = (2625 + 925) ÷ 100 = 3550 ÷ 100 = 35.5

That is exactly why the periodic table lists chlorine as 35.5 rather than a whole number, it is an average over isotopes, not the mass of any single atom. If a question gives you the percentages and isotope masses, the method is always this weighted average; if it gives you the relative atomic mass and one percentage, you rearrange the same equation to find the other.

Why this chapter pays off

Getting atomic structure fluent makes the next several chapters much easier, because you stop re-deriving the basics under exam pressure. A few habits help: always work from proton and nucleon numbers rather than guessing, write electron arrangements out fully, and be precise about the isotope definition, “same proton number, different nucleon number” earns the mark where “different atoms of an element” does not. Pinning down these exact phrasings with a chemistry glossary is worthwhile, and working steadily through the Matter and the Atomic Structure chapter builds the foundation the rest of the syllabus depends on. If the numbers or the isotope idea still feel slippery, this is a topic where a short, focused explanation clears it up quickly, and once it clicks, it stays clicked.