Isotopes and ions: atomic number and mass number, how isotopes differ, how ions form, and the standard nuclear notation.

What this dot point is asking

AQA wants you to define atomic number and mass number, use the standard nuclear notation, explain what isotopes are, and explain how atoms gain or lose electrons to form ions. This is part of topic 4.4.1 of the AQA GCSE Physics (8463) specification and overlaps directly with atomic structure in GCSE Chemistry.

Atomic number and mass number

The atomic number is the most important single fact about an atom because it fixes the identity of the element. Change the number of protons and you have a different element entirely. A neutral atom always has the same number of electrons as protons, so the atomic number also tells you the electron count and therefore the chemistry. The mass number, by contrast, tells you the total count of heavy particles (nucleons) in the nucleus, so it sets the mass of the atom. Protons and neutrons each have a relative mass of about $1$, while an electron has a relative mass of only about $1/1835$, so electrons make a negligible contribution to the mass number.

Nuclear notation

An element is written with the mass number on top and the atomic number below, next to the chemical symbol. For example, carbon written as mass number $12$ and atomic number $6$ has $6$ protons and $12 - 6 = 6$ neutrons. AQA may give you this notation and ask you to extract the particle counts, or describe an atom in words and ask you to deduce them. Always start from the atomic number for protons, match electrons to protons for a neutral atom, then subtract to find neutrons.

Isotopes

Because they have the same number of electrons, isotopes have the same chemical behaviour, but their nuclei may behave differently. Some isotopes are stable, while others are radioactive because their particular balance of protons and neutrons is unstable. Carbon has three naturally occurring isotopes: carbon-12 (the common stable form), carbon-13 (also stable), and carbon-14 (radioactive, used in carbon dating). All three have $6$ protons and so all behave identically in chemical reactions such as photosynthesis, which is why living things take in carbon-14 in the same proportion as it exists in the atmosphere. The differing neutron count is invisible to chemistry but decisive for nuclear stability.

The existence of isotopes is also why relative atomic masses on the periodic table are rarely whole numbers. The relative atomic mass is the average mass of the isotopes of an element, weighted by how common each one is. Chlorine, for instance, has a relative atomic mass close to $35.5$ because it is a mixture of roughly three parts chlorine-35 to one part chlorine-37.

Ions

It is electrons that are gained or lost, never protons or neutrons, because electrons sit on the outside of the atom and are held far more weakly than the particles bound inside the nucleus. A sodium atom with $11$ protons and $11$ electrons can lose one electron to become a sodium ion with $11$ protons but only $10$ electrons, giving it a net charge of $+1$. It is still sodium, because the proton count has not changed. An oxygen atom can gain two electrons to become an oxygen ion with a charge of $-2$. Ions are central to how electricity passes through solutions and molten compounds, and to how charge builds up in static electricity, where electrons are physically transferred from one surface to another by rubbing.

Try this

Q1. Define the term isotope. [2 marks]

• Cue. Atoms of the same element with the same number of protons but different numbers of neutrons.

Q2. An atom has atomic number $17$ and mass number $35$. State the number of protons, neutrons and electrons in the neutral atom. [3 marks]

• Cue. $17$ protons, $35 - 17 = 18$ neutrons, and $17$ electrons.