Atomic structure: protons, neutrons and electrons, atomic number and mass number, isotopes and relative atomic mass, and the development of the model of the atom.

What this dot point is asking

Edexcel wants you to describe the atom in terms of protons, neutrons and electrons, give their relative masses and charges, use atomic number and mass number to work out the particles in any atom or ion, explain what isotopes are and calculate relative atomic mass from isotopic abundances, and outline how the model of the atom changed as new evidence appeared. This is the foundation of Paper 1 and feeds straight into bonding and the calculations.

The subatomic particles

Every atom is built from three particles. Their relative masses and charges are worth memorising exactly.

The nucleus is positively charged (protons and neutrons) and contains almost all the mass, because electrons are about $1836$ times lighter than a proton. The atom is mostly empty space: if the atom were the size of a stadium, the nucleus would be a pea in the centre. The typical radius of an atom is about $1 \times 10^{-10}$ m, while the nucleus is roughly $1 \times 10^{-14}$ m across, about ten thousand times smaller.

Atomic number and mass number

Two numbers describe any atom, written in the form $^{A}_{Z}X$:

• Atomic number $Z$ is the number of protons. It defines the element, and in a neutral atom it also equals the number of electrons.
• Mass number $A$ is the total number of protons and neutrons.

So the number of neutrons is $A - Z$, the mass number minus the atomic number. For example, sodium $^{23}_{11}Na$ has $11$ protons, $11$ electrons and $23 - 11 = 12$ neutrons.

Isotopes

Because isotopes have the same number of electrons and the same electronic configuration, they have identical chemical properties. They differ only slightly in physical properties such as density, because their masses differ. Carbon, for instance, exists as $^{12}C$, $^{13}C$ and $^{14}C$, all with $6$ protons but $6$, $7$ and $8$ neutrons respectively.

Relative atomic mass

Because an element is usually a mixture of isotopes, we use a weighted average called the relative atomic mass ($A_r$).

The development of the model of the atom

The model changed as experiments produced new evidence. You should be able to say what each scientist contributed and what evidence drove the change.

1. John Dalton (early 1800s) proposed that elements are made of tiny indivisible solid spheres, and that atoms of different elements differ.
2. J. J. Thomson (1897) discovered the electron and proposed the plum pudding model: a ball of positive charge with negative electrons embedded in it.
3. Ernest Rutherford (1911) carried out the alpha-particle scattering experiment: most alpha particles passed straight through gold foil, but a few were deflected back. This showed the atom is mostly empty space with a tiny, dense, positively charged nucleus, replacing the plum pudding model.
4. Niels Bohr refined the model by placing electrons in fixed shells (energy levels) at set distances from the nucleus, which explained why electrons are not pulled into the nucleus.
5. Later work identified the proton in the nucleus, and James Chadwick (1932) discovered the neutron, completing the modern nuclear model.

Try this

Q1. State the relative charge and relative mass of a neutron. [2 marks]

• Cue. Relative charge $0$; relative mass $1$.

Q2. An atom has $17$ protons and $18$ neutrons. Give its atomic number and mass number. [2 marks]

• Cue. Atomic number $= 17$; mass number $= 17 + 18 = 35$.

Q3. Copper has two isotopes, $^{63}Cu$ ($69\%$) and $^{65}Cu$ ($31\%$). Calculate the relative atomic mass of copper to one decimal place. [3 marks]

• Cue. $[(63 \times 69) + (65 \times 31)] / 100 = (4347 + 2015) / 100 = 6362 / 100 = 63.6$.