It makes an impression to think that for two thousand years mankind thought that the world was made of small, indivisible particles called atoms. It makes an impression to think that this concept was a figment of the imagination of a philosopher, Democritus in the 5th century BC, without scientific experimental support, and it makes an impression that this theory was the basis of modern science until the early 20th century.

This atomic model began to falter with the discovery of radioactivity and the discovery of a negative particle called the electron.

These discoveries made it clear that if the atom can lose particles during radioactive decay, and

there are electrons inside it, then it is not indivisible, it is not the smallest part of the matter.

The discovery of the electron led to the hypothesis of new atomic models, such as that of the British physicist Joseph John Thomson who, after discovering the electron, hypothesised that the atom was composed of a positive cloud in which negative corpuscles were present. Thomson’s model went down in history as the ‘sultana loaf model’.

This model lasted for 14 years, nothing like Democritus’ two thousand, until the coronation of Rutherford’s model, who realised that the atom was essentially empty except for a central, denser area called the nucleus. This nucleus was composed of protons and neutrons, and electrons revolve around it. This model was called the ‘planetary model’ because of its similarity to our Solar System.

This atomic model seemed to be reliable, given the experimental evidence at the time, but it was not. The model fell apart with the arrival of Bohr, who explored the subject of electron orbits around the nucleus, realising that electrons could not simply orbit the nucleus, but had to do so in stable orbits.

To understand this, imagine we are in space, in a space station orbiting a planet. Our orbit is nothing special; if we turn on the rocket propulsion, our space station can move 7 meters away from the planet and turn onto a new orbit, descend a few meters and follow another orbit.

In short, all orbits are permitted, but this is not the case for electrons.

The orbits they can travel are fixed at defined distances from the nucleus. The orbits according to Bohr’s model are like steps on a ladder, the electron cannot stand halfway between two steps, otherwise, it would risk falling back onto the previous one.

From Bohr’s model of 1913, we move on to Schrodinger’s model of 1926, which is the model

currently in use to describe the modern atom.

According to Schrodinger’s model, electrons move around the nucleus on fixed orbitals; we do not know the exact location of the electrons, but we only know that they have a certain probability of being found in certain regions of space surrounding the atom, which are called orbitals. Orbitals are a region of space where the probability of finding an electron is highest, at least 90%.

Thanks to this model, we now know how to represent the electrons in orbitals around the nucleus and are thus able to predict their behaviour when interacting with other atoms. The electrons are arranged in the orbitals according to very precise rules, starting with the first rule that places them so that the atom’s total energy is minimal, attractions are maximised and repulsions are minimised, by Pauli’s exclusion principle, which states that each orbital can contain a maximum of two electrons; and Hund’s rule, which tells us that when orbitals with the same energy are to be filled, an electron is placed on each orbital and then the half-filled orbitals are filled. More orbitals can exist at the same energy.

Representing the electrons around the nucleus means assigning the fundamental-state configuration to a chemical element of atomic number Z, which corresponds to the number of protons in the nucleus of an atom (in a neutral atom, the number of protons is equal to the number of electrons); to represent the electronic configuration of the fundamental state, we apply the following criteria:

We add Z electrons, one after the other, associating no more than two electrons with each orbital (Pauli exclusion principle).

If there is more than one orbital in a given sub-level, we add electrons to distinct orbitals in the sub-level under consideration before completely occupying one orbital (Hund’s rule). We know that for each energy level, there can be more than one orbital.

We identify the orbitals by letters of the alphabet according to increasing energy, with a superscript indicating the number of electrons associated with the orbital under consideration. The configuration of a complete level is represented by the symbol of the noble gas with that configuration, as in [He] for 1s2.

The procedure described provides the electronic configuration of the fundamental state; any other arrangement will correspond to an excited state. Note that the structure of the periodic table makes it possible to predict the electronic configuration of most elements once it is clear which orbitals are filled in each block of the table.

When is an electron in an excited orbital state? When it is in an orbital that is not in the fundamental state, i.e. not in a minimum energy position. And this may be because it has captured an energy packet (photon) that has excited it, but it can return to its initial state by emitting a new photon (e.g. phosphorescent minerals).