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Module 2: Foundations in chemistry›Atomic structure and isotopes (2.1.1)

Atomic structure and isotopes (2.1.1)

The structure of the atom, the existence of isotopes, relative mass, and the use of mass spectroscopy for determining isotopic abundance.
3 min

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The structure of the atom:

A diagram illustrating an atom, featuring a blue nucleus at the center and several yellow electrons orbiting around it in concentric circles.

The nucleus contains the protons and neutrons.

The nucleus is surrounded by shells (orbitals) which contain the electrons.

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Subatomic particles can be described by their relative charge and relative mass.

A table displaying information about subatomic particles, including their names, relative charges, and relative masses. The table lists protons with a charge of +1 and mass of 1, neutrons with a charge of 0 and mass of 1, and electrons with a charge of -1 and mass of 1/1837.
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In a neutral atom, the number of protons is the same as the number of electrons.

Atomic number () = number of protons
Mass number () = number of neutrons + the atomic number

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The number of neutrons = mass number − the atomic number.

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Ions are particles with a charge.

In an ion, the number of electrons = number of protons − the charge.

  • The number of electrons in a
  • The number of electrons in a
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Dalton (1800s):

  • discovered that atoms can be combined to make compounds
  • believed atoms cannot be broken into smaller parts
  • disproved by J.J. Thomson’s cathode ray tube experiments.
A simple illustration of a light blue sphere on a white background, with a thin black outline.
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Thomson (1897):

  • discovered the electron
  • believed that atoms were a sphere of positive mass embedded with electrons: the plum pudding model
  • was disproved by Rutherford’s gold foil experiment.
Illustration of the 'plum pudding' model of the atom, showing a positively charged sphere with negatively charged electrons represented as small circles within it.
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Rutherford (1905):

  • discovered the nucleus with his gold foil experiment.
Illustration of an atom showing a blue nucleus at the center with three yellow electrons orbiting around it, labeled 'Nucleus' and 'Electron'.
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Rutherford’s gold foil experiment:

A diagram illustrating the interaction of alpha particles with atoms. The image shows several atoms represented as circles with a blue center and a yellow outer layer. Arrows indicate the paths of alpha particles approaching and interacting with the atoms, labeled as (A), (B), and (C).
  • A very thin piece of gold foil was bombarded with alpha particles.
  • Most particles went straight through (A).
  • Some were deflected (B) and a few were reflected (C).
  • This proved that the vast majority of the mass of an atom is concentrated in a small central volume and that the central mass was positively charged.
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Bohr (1913):

  • discovered energy levels
  • related atomic spectra to electronic energy levels.
A diagram illustrating an atom, featuring a blue nucleus at the center and several yellow electrons orbiting in circular paths around it.
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Isotopes are atoms of the same element with different numbers of neutrons and different masses.

A comparison of the atomic structures of Carbon-12 and Carbon-13. On the left, Carbon-12 is shown with 6 neutrons (red spheres) and 6 protons (purple spheres), with an atomic mass of 12.000. On the right, Carbon-13 has 7 neutrons and 6 protons, with an atomic mass of 13.003. Both atoms have 6 electrons (yellow spheres) orbiting around the nucleus.

Carbon-12 and carbon-13 are isotopes of carbon.

  • Both identify as carbon, with an atomic number of six.
  • Carbon-12 has six neutrons and a relative atomic mass of 12 whereas carbon-13 has seven neutrons and a relative atomic mass of 13.
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Relative isotopic mass: mass of a specific isotope compared with the mass of carbon-12.

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Relative atomic mass: weighted mean mass of all isotopes in an element, compared with the mass of carbon-12.

It is determined using the following equation:

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The term relative formula mass (RFM) is used for compounds with giant structures like ionic compounds or giant covalent lattices.

The RFM and (relative molecular mass) can be found from the sum of the relative atomic masses, .

Taking as an example:

  • Sodium has
  • Sulfur has
  • Oxygen has

RFM

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The mass spectrum below shows the relative abundance of the two isotopes of boron.

A bar graph displaying relative abundance on the y-axis and m/z values on the x-axis. The graph shows a significant peak at m/z 11 with a relative abundance of 100, a smaller peak at m/z 10 with a relative abundance of 23, and no other notable peaks.

Isotope 1: m/z , relative abundance
Isotope 2: m/z , relative abundance

The relative atomic mass of boron can be calculated from this data:

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