Edexcel GCSE Chemistry

How the model of the atom changed

At the start of the nineteenth century John Dalton described atoms as tiny solid spheres that could not be divided, with a different kind of sphere for each element. That model explained why elements combine in fixed proportions, but it had no room for anything smaller than an atom, so it had to change when smaller particles were found.

In 1897 J. J. Thomson showed that atoms contain negatively charged particles, later called electrons, which are far lighter than any atom. Atoms are neutral overall, so he pictured a ball of positive charge with electrons embedded in it, like fruit in a pudding. In 1909 Geiger and Marsden, working with Rutherford, fired positively charged alpha particles at thin gold foil. Most went straight through, a few were deflected through large angles and a very few bounced almost straight back. A ball of spread-out positive charge could not do that, so Rutherford proposed that the positive charge and nearly all the mass are packed into a tiny nucleus at the centre, with the electrons around it and mostly empty space in between.

Niels Bohr then suggested that the electrons move in fixed orbits, or shells, at set distances from the nucleus, which fitted observations of the light that atoms give out. Later work showed that the nucleus itself contains positively charged protons, and in 1932 James Chadwick discovered the neutron, a particle with no charge and about the same mass as a proton. That completed the model taught today: a nucleus of protons and neutrons, surrounded by electrons in shells.

Protons, neutrons and electrons

Atoms have no overall charge, so the number of protons is always equal to the number of electrons: the positive charge of the nucleus is cancelled exactly by the negative charge of the electrons around it.

The nucleus is very small compared with the whole atom. An atom has a radius of about $1 \times 10^{-10}$ m, and the nucleus is roughly ten thousand times smaller than that, about $1 \times 10^{-14}$ m. If the nucleus were the size of a pea, the atom would be the size of a large sports stadium. Electrons have almost no mass, so the nucleus holds nearly all of an atom's mass.

Atomic number, mass number and isotopes

The atomic number is the number of protons in the nucleus. It is unique to each element: every atom with 6 protons is carbon, and no other element has atoms with 6 protons. The atomic number also fixes an element's place in the periodic table.

The mass number is the total number of protons and neutrons in the nucleus. Both are written beside the symbol, with the mass number at the top:

This atom has 17 protons, since the atomic number is 17, and therefore 17 electrons. The number of neutrons is the mass number minus the atomic number, $35 - 17 = 18$.

Isotopes are atoms of one element whose nuclei hold the same count of protons and different counts of neutrons, so they share an atomic number and differ in mass number. Carbon-12 and carbon-14 are both carbon, with 6 protons each, but carbon-12 has 6 neutrons and carbon-14 has 8. Isotopes of an element have the same chemical properties, because chemical behaviour depends on the electrons, and the electron arrangement is decided by the number of protons.

Relative atomic mass

The relative atomic mass, $A_r$, of an element is the average mass of its atoms compared with carbon-12 (which is given the value 12). It is not usually a whole number, because most elements exist as a mixture of isotopes. Chlorine is an example: it has an $A_r$ of 35.5 because a sample contains about three atoms of chlorine-35 for every one of chlorine-37, so the average lies between the two mass numbers, closer to 35.

Higher tier only: the relative atomic mass can be calculated from the relative masses of the isotopes and how much of each there is (their abundance). Multiply each isotope's mass number by its percentage abundance, add the results and divide by 100.

Boron has two isotopes: boron-10 is 20% of the atoms and boron-11 is 80%.

Magnesium has three isotopes, of mass 24, 25 and 26, with abundances of 79%, 10% and 11%.

A quick check is that the answer must lie between the lightest and heaviest isotope, and closer to the most abundant one. Both answers do.