Constituents of the Atom
The Bohr-style picture of an atom — a small dense nucleus with electrons around it — is the starting point for the whole of particle physics. This lesson sets out the three particles that model is built from, the notation used to label a given nucleus, and what it means for two atoms to be isotopes of each other.
The simple model of the atom
An atom is modelled as a tiny, dense, positively charged nucleus at its centre, made of protons and neutrons, surrounded by a cloud of electrons occupying a much larger volume of essentially empty space. Protons and neutrons are collectively called nucleons. Almost all of an atom's mass sits in the nucleus, while almost all of its volume is taken up by the electrons orbiting far outside it.
Each particle has a definite charge and mass, worth knowing in two forms: the SI values (coulombs, kilograms), and the simpler relative values, where the proton's charge and mass are each taken as the unit of comparison.
| Particle | Charge / C | Relative charge | Mass / kg | Relative mass |
|---|---|---|---|---|
| Proton | $+1.60\times10^{-19}$ | $+1$ | $1.673\times10^{-27}$ | $1$ |
| Neutron | $0$ | $0$ | $1.675\times10^{-27}$ | $\approx 1$ |
| Electron | $-1.60\times10^{-19}$ | $-1$ | $9.11\times10^{-31}$ | $\approx \tfrac{1}{1836}$ |
The proton and neutron have almost identical mass, close enough that both are treated as having a relative mass of exactly 1, while the electron is about 1836 times less massive. This gap is why the electron dominates chemistry and everyday electrical behaviour, while the nucleus dominates the mass.
Every charge you'll meet at this level, on any particle, comes in whole-number multiples of the elementary charge $e = 1.60\times10^{-19}\,\text{C}$. That's why it's so convenient to quote charges as simple relative numbers like $+1$, $-1$ or $0$ rather than repeating the same $10^{-19}$ factor every time.
Specific charge
The specific charge of a particle is its charge-to-mass ratio:
measured in coulombs per kilogram (C kg⁻¹). Some specifications call this the charge-to-mass ratio — the same quantity under a different name.
Putting actual numbers through that ratio gives the specific charge in coulombs per kilogram directly. For the proton:
and for the electron, with the same size of charge but the much smaller mass:
Checking the ratio of these two numbers confirms the reasoning above: $1.76\times10^{11} / 9.58\times10^{7} \approx 1836$, exactly the mass ratio expected.
Specific charge is large when a particle is light for the amount of charge it carries, and small when it's heavy for its charge. A neutron, with no charge, has a specific charge of exactly zero.
Proton number, nucleon number, and nuclide notation
Every nucleus is fully specified by two whole numbers:
- The proton number (or atomic number) $Z$ — the number of protons, which fixes which element the atom is and the number of electrons in a neutral atom.
- The nucleon number (or mass number) $A$ — the total number of protons and neutrons, $A = Z + N$, where $N$ is the number of neutrons.
These are written alongside the element's chemical symbol X in nuclide notation:
$A$ is written as a superscript before the symbol, $Z$ as a subscript directly below it. For example, $^{12}_{6}\text{C}$ denotes a carbon nucleus ($Z=6$) with nucleon number 12, containing $12-6=6$ neutrons.
Top number ($A$) = protons + neutrons. Bottom number ($Z$) = protons only. The number of neutrons is never written directly — you always find it by subtracting: $N = A - Z$. Because the chemical symbol already tells you the element (and therefore $Z$), the subscript is sometimes dropped in informal writing, but you should always include it in exam answers unless told otherwise.
Isotopes
Atoms with the same proton number $Z$ but a different nucleon number $A$ (the same number of protons but a different number of neutrons) are called isotopes of each other.
Since isotopes share the same $Z$, they have the same number of electrons arranged the same way, so they behave almost identically in chemical reactions — carbon-12 and carbon-14 form the same compounds.
What differs is anything depending on the nucleus itself: mass, and nuclear stability. Extra neutrons add mass without adding charge, so isotopes have measurably different masses, and some combinations of protons and neutrons form stable nuclei while others are radioactive.
Key takeaways
- The atom is modelled as a small, dense, positive nucleus of protons and neutrons (nucleons), surrounded by orbiting electrons.
- Proton: charge $+1.60\times10^{-19}\,\text{C}$ (relative $+1$), mass $1.673\times10^{-27}\,\text{kg}$ (relative $1$).
- Neutron: charge $0$, mass $1.675\times10^{-27}\,\text{kg}$ (relative $\approx 1$).
- Electron: charge $-1.60\times10^{-19}\,\text{C}$ (relative $-1$), mass $9.11\times10^{-31}\,\text{kg}$ (relative $\approx 1/1836$).
- Specific charge $= Q/m$: proton $= 9.58\times10^{7}\,\text{C}\,\text{kg}^{-1}$, electron $= 1.76\times10^{11}\,\text{C}\,\text{kg}^{-1}$ — because the proton and electron carry equal-sized charge but very different mass, the electron's specific charge is about 1836 times larger in magnitude.
- Nuclide notation $^{A}_{Z}\text{X}$: $A$ (top) is the nucleon number (protons + neutrons), $Z$ (bottom) is the proton number (protons only); $N = A - Z$ gives the neutron number.
- Isotopes share the same $Z$ but have different $A$ — same chemistry (same electron arrangement), different mass and different nuclear stability.