HSC Chemistry

Subatomic Particles, Isotopes and Relative Atomic Mass: HSC Chemistry Explained

Published 2026-10-12 · Updated 2026-10-12

Relative atomic mass is one of those topics where students lose marks not because the chemistry is hard, but because they confuse mass number with atomic mass, or forget that the periodic table value is a weighted average rather than a whole number. HSC markers consistently see students treat chlorine's atomic mass of 35.5 as an error rather than a deliberate, calculated result. Getting the fundamentals right here also unlocks isotope questions, ion questions, and nuclear stability, so it is worth spending the time to build a solid picture from the ground up.

What the syllabus asks

This concept sits within Module 1: Properties and Structure of Matter in the NSW HSC Chemistry course. Students need to understand the structure of the modern atomic model, the properties of subatomic particles, how atomic number and mass number are used, what isotopes are, and how relative atomic mass is calculated from isotopic abundances. The concept also extends to the stability of isotopes and how the neutron-to-proton ratio determines whether a nucleus is stable or radioactive.

The idea, explained

The modern atomic model

Atoms consist of a small, dense nucleus containing protons and neutrons, surrounded by an electron cloud containing electrons. The nucleus holds almost all of the atom's mass, while the electron cloud determines how the atom behaves chemically.

Subatomic particles

There are three subatomic particles you need to know:

• Proton: relative charge +1, relative mass 1, located in the nucleus.
• Neutron: relative charge 0, relative mass 1, located in the nucleus.
• Electron: relative charge -1, relative mass approximately 1/1836, located in the electron cloud.

Because electrons have such a tiny mass compared to protons and neutrons, they contribute negligibly to the overall mass of an atom.

Atomic number, mass number and neutrons

The atomic number (Z) is the number of protons in the nucleus. It defines which element an atom is. The mass number (A) is the total number of protons and neutrons (collectively called nucleons). From these two values, the number of neutrons is simply A minus Z.

In nuclide symbol notation, the mass number is written at the top-left and the atomic number at the bottom-left of the element symbol. For example, in ¹⁴₆C, Z = 6 (protons) and A = 14, so the number of neutrons = 14 - 6 = 8.

Neutral atoms and ions

In a neutral atom, the number of electrons equals the number of protons. When electrons are lost, the atom becomes a positively charged ion called a cation. When electrons are gained, the atom becomes a negatively charged ion called an anion. Importantly, gaining or losing electrons does not change the atomic number or mass number.

Isotopes

Isotopes are atoms of the same element that have different numbers of neutrons. Because they share the same atomic number, they occupy the same position in the periodic table (same period, same group). Carbon-12, carbon-13 and carbon-14 are all in period 2, group 14. Because they have the same number of electrons, isotopes have the same chemical properties. Their physical properties differ slightly due to the difference in mass.

Carbon-12 (¹²₆C) and carbon-13 (¹³₆C) are stable isotopes. Carbon-14 (¹⁴₆C) is unstable because it has too many neutrons relative to its protons, and it undergoes beta-minus decay.

Stable and unstable isotopes

Whether an isotope is stable depends on its neutron-to-proton (n/p) ratio. Nuclei whose n/p ratio falls within the band of stability are stable and do not decay. For lighter elements (Z less than 20), stability occurs when the n/p ratio is approximately 1, meaning roughly equal numbers of neutrons and protons. For heavier elements, a higher n/p ratio is needed to overcome the repulsion between protons.

Isotopes outside the band of stability are radioactive and undergo spontaneous decay. If there are too many neutrons relative to protons, beta-minus decay occurs. If there are too many protons relative to neutrons, beta-plus decay or electron capture occurs. Very heavy nuclei with large atomic numbers tend to undergo alpha decay.

Relative atomic mass

Because most elements exist as a mixture of isotopes in nature, the atomic mass shown on the periodic table is not a whole number. It is a weighted average of the masses of all naturally occurring isotopes, taking their fractional abundances into account.

The formula is: Ar = sum of (isotope mass multiplied by fractional abundance)

For chlorine, which has two naturally occurring isotopes, Cl-35 (75% abundance) and Cl-37 (25% abundance):

Ar = (35 × 0.75) + (37 × 0.25) = 26.25 + 9.25 = 35.5

This is exactly why chlorine's atomic mass on the periodic table is 35.5 rather than a whole number.

What the exam asks

HSC examination questions on this topic typically ask students to:

• Identify the number of protons, neutrons and electrons in a neutral atom or ion given a nuclide symbol.
• Explain what isotopes are and describe how they differ from one another.
• Calculate relative atomic mass from given isotopic masses and percentage abundances.
• Explain why atomic masses on the periodic table are not whole numbers.
• Assess whether an isotope is likely to be stable or unstable based on its n/p ratio.

For calculation questions, always convert percentage abundances to decimal fractions before substituting into the formula.

Common mistakes

• Confusing mass number with relative atomic mass. Mass number is always a whole number (it counts nucleons). Relative atomic mass is a weighted average and is usually a decimal.
• Forgetting to convert percentages to fractions. Using 75 instead of 0.75 in the Ar formula gives an answer 100 times too large.
• Thinking isotopes have different chemical properties. Isotopes have the same number of electrons, so their chemical behaviour is the same. Only physical properties differ slightly.
• Assuming a neutral atom always has equal protons and electrons after ionisation. Once an atom gains or loses electrons to form an ion, the electron count changes, but the proton count (and therefore the element identity) does not.
• Placing isotopes in different positions on the periodic table. Isotopes of the same element always occupy the same position because position is determined solely by atomic number.

If you want to work through practice questions and get instant feedback on your working, Avocado's AI-native tutoring service has a dedicated lesson on Subatomic Particles, Isotopes and Relative Atomic Mass that walks you through every part of this concept at your own pace.