Although it’s often convenient to think of atoms as hard spheres (like marbles), the reality is a little different. Atoms are composed of a dense, positively charged nucleus (made up of positively charged protons, p+ and neutral neutrons, n) surrounded by a “cloud” of negatively charged electrons, e–. There are equal numbers of protons and electrons, so that the positive and negative charges cancel each other out to give a neutral atom overall. When atoms gain or lose electrons and become charged (atoms cannot lose protons or neutrons under normal conditions), they are called ions.
The electron cloud is not composed of electrons directly, but actually represents a probability function. The denser an area of the cloud is, the more likely we are to find an electron there. This is a consequence of wave-particle duality (a central concept in quantum mechanics, which states that electrons and other atomic-scale objects show both wave-like and particle-like properties) and the Heisenberg uncertainty principle (another central concept, which states that you cannot know both the position and the momentum of a particle a the same time).
Early in the 20th century, more advanced experimental apparatus led to several new theories about atomic structure being developed. Probably the most well-known is the Rutherford model, which has electrons orbiting a small, dense nucleus much like planets orbit the sun. The straightforward aesthetics of this image have led to it being widely adopted as a symbol of nuclear power (for example, in the logos of the IAEA and the Albuquerque Isotopes).
When full-blown theories of quantum mechanics were adopted, though, the models had to change. We currently understand the electron cloud in terms of atomic orbitals – specific regions where certain electrons are likely to be found. Each orbital represents a different probability density function.
If an atom is very large, it will have lots of protons in its nucleus. It therefore needs lots of electrons to balance all the positive charge, and hence needs lots of atomic orbitals to accommodate them. Conversely, a small atom may only need one or two orbitals.
Atomic orbitals are described by three quantum numbers:
- The principal quantum number n tells us about orbital energy (referred to as an energy level)
- The orbital angular momentum quantum number l tells us about orbital shape
- The magnetic quantum number ml tells us about orbital alignment
For example, an s-orbital (l=0) is spherical, and so is the same in all directions. A p-orbital (l=1) is figure-of-eight shaped, and may be aligned in one of three directions (commonly represented by the three Cartesian axes x, y and z).
When atoms form bonds to make molecules, they are sharing electrons and the atomic orbitals overlap to form molecular orbitals. These are quite similar to the atomic orbitals, but the electrons are delocalised – spread out over the entire molecule rather than being concentrated on the individual atoms. This means that when we study the electrons in a molecule, we cannot tell which atom the electrons came from – they “belong” to the molecule as a whole. Molecular orbitals are much more complicated than atomic orbitals, as there are several different nuclei and many more electrons for us to consider, but share many of the same properties.