Each particle has its own mass and charge.

A proton has a relative mass of 1 and a relative charge of +1.

A neutron has a relative mass of 1 and has no charge.

An electron has a relative charge of −1, but a mass so small (~1/1836 of a proton) that it is ignored.

This means the mass of an atom is almost entirely concentrated in its nucleus.

Every relative atomic mass is defined against the same benchmark: 1/12 of the mass of an atom of Carbon-12.

In a neutral atom, the number of positively charged protons is balanced by an equal number of negatively charged electrons, so the overall charge is zero.

If an atom gains or loses electrons, it becomes an ion. A positive ion (cation) has lost electrons, a negative ion (anion) has gained them.

The number of protons in an atom’s nucleus is its atomic number, given the symbol Z.

Z is what defines an element. Any hydrogen has 1 proton, any carbon has 6, any oxygen has 8, and so on.

Change the number of protons, and you’ve changed the element entirely.

The mass number, A, is the total number of protons and neutrons in the nucleus.

For carbon-12, the mass number is 12 - six protons and six neutrons.

For carbon-14, the mass number is 14 - still six protons, but eight neutrons.

These two atoms are both carbon, but they’re different isotopes.

Isotopes are atoms of the same element with the same number of protons but a different number of neutrons.

Because chemical reactions involve the electrons of an atom, not the nucleus, all isotopes of an element behave the same way chemically.

They do differ slightly in physical properties (density, rate of diffusion) because these depend on mass.

Every relative atomic mass is defined against the same benchmark: 1/12 of the mass of an atom of carbon-12.