IB Chemistry SLTopic 1 — The Nuclear AtomPaper 1 & 2Core idea~9 min read
Isotopes
Not every atom of an element is identical. Isotopes are atoms of the same element with the same protons but a different number of neutrons — same chemistry, slightly different mass. That small difference is exactly what lets us calculate an element’s relative atomic mass.
📘 What you need to know
Isotopes are atoms of the same element with the same number of protons (and electrons) but a different number of neutrons.
They therefore have the same atomic number but different mass numbers.
Isotopes have the same chemical properties (same electron arrangement) but different physical properties (different mass).
Relative atomic mass (Ar) is the average mass of an element’s atoms compared to 1/12 of a carbon-12 atom.
Ar is worked out from the percentage abundance of each isotope.
What is an isotope?
Because the number of protons defines the element, changing the neutrons doesn’t make a new element — it makes a different isotope of the same one.
Carbon-12 and carbon-14 are both carbon: each has 6 protons, but they have 6 and 8 neutrons.
You’ll see them written as 12C / C-12 and 14C / C-14.
The three isotopes of hydrogen: same single proton, different numbers of neutrons.
Same chemistry, different physics
Isotopes behave identically in reactions but differ in mass — and that split comes down to which particles change.
Same chemical properties: isotopes have the same number of electrons in the same arrangement, and chemistry is all about electrons.
Different physical properties: the extra neutrons change the mass, so things like density, melting/boiling point and rate of diffusion differ slightly.
Quick way to remember it: reactions are decided by electrons, and isotopes have identical electrons — so identical chemistry. It’s only the neutrons (hence mass) that differ.
Relative atomic mass
Because an element is usually a mix of isotopes, we quote an average mass — the relative atomic mass (Ar). Formally it’s the average mass of an atom compared with 1/12 of a carbon-12 atom.
We work it out by weighting each isotope’s mass by how common it is (its percentage abundance):
Relative atomic mass
Ar = [ (% × mass) + (% × mass) + … ] ÷ 100
WORKED EXAMPLE
A sample of oxygen contains 16O (99.76%), 17O (0.04%) and 18O (0.20%). Calculate the relative atomic mass of oxygen to 2 decimal places.
Weight each isotope by its abundance(99.76 × 16) + (0.04 × 17) + (0.20 × 18)= 1596.16 + 0.68 + 3.60 = 1600.44Divide by 1001600.44 ÷ 100 = 16.0044A_r = 16.00
💡 Exam tip
Only round at the very end — carry the full figure through the calculation, then round to the decimal places asked for.
The percentage abundances can be given or read off a mass spectrum; the method is the same either way.
That completes The Nuclear Atom section. Next comes Electronic Configurations — how those electrons are arranged in energy levels and orbitals around the nucleus.
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