Edexcel A-Level Chemistry: Atomic Structure and Mass Spectrometry — Revision Notes

Condensed for the final weeks. For the full explanation, use the Atomic Structure and Mass Spectrometry study guide.

Subatomic particles (2.1–2.4)

Particle	Relative mass	Relative charge
Proton	1	+1
Neutron	1	0
Electron	Negligible	−1

Atomic (proton) number = number of protons = number of electrons in a neutral atom. Mass number = protons + neutrons. Neutrons = mass number − atomic number.

For an ion, adjust the electron count by the charge: a positive ion has fewer electrons than protons; a negative ion has more electrons than protons. Practise this subtraction/addition until automatic — it is tested constantly.

Isotopes (2.5)

Isotopes = atoms of the same element (identical proton number) with different numbers of neutrons, and therefore different mass numbers. Because chemistry is governed by electron arrangement, isotopes of the same element have (near-)identical chemical properties but different physical properties (e.g. mass).

Mass spectrometry (2.6–2.7)

A mass spectrometer ionises a sample and separates the resulting ions by mass-to-charge ratio, producing a mass spectrum used to:

• Deduce a sample’s isotopic composition.
• Calculate relative atomic mass from isotope abundances.
• Determine relative molecular mass from a molecular ion peak, to help identify a compound.
• Recognise that ions may carry a 2+ charge, not just 1+.

Relative atomic mass = weighted mean, not a simple average:

Ar = (mass1 x abundance1 + mass2 x abundance2 + ...) / 100

Worked example (chlorine): ³⁵Cl (75%) and ³⁷Cl (25%).

35 x 75 = 2625
37 x 25 = 925
(2625 + 925) / 100 = 35.5

Ar = 35.5 — matches chlorine’s known value, and shows why relative atomic masses are rarely whole numbers.

Predicting mass spectra for diatomic molecules (2.7)

Chlorine is diatomic (Cl₂), so its molecular ion peaks reflect combinations of its two isotopes — not simply one peak per isotope. Two isotopes combine into three possible molecular ion masses:

Combination	Mass
³⁵Cl–³⁵Cl	70
³⁵Cl–³⁷Cl	72
³⁷Cl–³⁷Cl	74

Relative peak heights follow the probability of each combination occurring, calculated from the isotopic abundances (not simply equal heights) — this is a common source of lost marks when students assume three equal peaks.

Worked example: identifying a molecule from its molecular ion peak

A mass spectrum shows a molecular ion peak at m/z = 44, for an unknown organic compound believed to contain only carbon, hydrogen and oxygen.

Step 1: list candidate molecular formulae with relative molecular
        mass 44
        - CO2 (carbon dioxide): 12 + 16 + 16 = 44
        - C3H8 (propane): (12 x 3) + (1 x 8) = 44
        - C2H4O (e.g. acetaldehyde): (12 x 2) + (1 x 4) + 16 = 44

Step 2: use other evidence (fragmentation pattern, the sample's
        known context, or other spectroscopic data) to distinguish
        between candidates sharing the same relative molecular mass

This illustrates why the molecular ion peak alone identifies relative molecular mass, but rarely identifies a compound uniquely without further evidence — a distinction worth stating explicitly in an exam answer rather than assuming one m/z value maps to one structure.

Why isotopic abundance data matters beyond one calculation

Relative atomic mass calculated from isotopic abundance is not a one-off exam trick — it is the reason the periodic table’s atomic masses are decimals rather than whole numbers for almost every element. Chlorine’s 35.5 reflects a natural mixture of ³⁵Cl and ³⁷Cl in the fixed ratio found in nature; a sample with an artificially altered isotope ratio (as used in some medical or industrial applications) would have a different relative atomic mass from the standard periodic table value, even though it is still chemically “chlorine.” Keeping this distinction between an element’s identity (fixed by proton number) and its measured relative atomic mass (dependent on the specific isotopic mixture present) clear is useful both for this outcome and for isotope-related content later in the specification.

Key terms

Atomic (proton) number — protons in an atom; equals electrons in a neutral atom; uniquely identifies the element. Mass number — protons + neutrons. Isotope — same proton number, different neutron number, different mass number. Relative atomic mass — weighted mean mass of an element’s atoms across its naturally occurring isotopes, on the carbon-12 scale. Mass spectrometer — ionises and separates ions by mass-to-charge ratio. Molecular ion peak — the peak corresponding to an entire, unfragmented molecule (or diatomic unit), used to find relative molecular mass.

Common mistakes

• Calculating relative atomic mass as a simple average instead of weighting by abundance.
• Confusing atomic number with mass number when finding the number of neutrons.
• Forgetting to adjust electron count for a charged ion.
• Assuming a diatomic element’s spectrum shows only as many peaks as isotopes (two for chlorine) rather than the full set of combinations (three).

Quick self-test

• State the relative mass and charge of a proton, neutron and electron.
• Find the number of protons, neutrons and electrons in a Mg²⁺ ion (atomic number 12, mass number 24).
• Calculate the relative atomic mass of an element with isotopes of mass 10 (20% abundance) and 11 (80% abundance).
• Explain why chlorine’s mass spectrum shows three molecular ion peaks, not two.
• Define isotope precisely, using proton and neutron number.

Related resources

Atomic Structure and Mass Spectrometry study guide | Atomic Structure and Mass Spectrometry practice questions

Official syllabus

Pearson Edexcel International Advanced Subsidiary/Advanced Level in Chemistry (XCH11/YCH11) specification, Issue 1, September 2017 — qualifications.pearson.com.