Nucleus Structure

Uncovering the atomic structure of matter was one of physics' most significant successes in the 20th Century, with Ernest Rutherford's famous gold foil experiment in the 1910s confirming the existence of the atomic nucleus. Rutherford's experiment involved firing positively charged alpha particles at atoms and measuring the angle of deflection. Initially, the results that they found baffled them and were totally at odds with the popular 'Plum Pudding' Model at the time. What Rutherford eventually realized was that their results could only be obtained if there was an extremely dense, positively charged region at the heart of the atom. This discovery kick-started the dawn of nuclear physics, a field of physics that would come to dominate the 20th Century and beyond. In this article, we'll look a little closer at the structure of the atom, in particular the nucleus and electrons. We'll also investigate how the different structures of atoms can give rise to different chemical and radioactive properties.

Atomic Nucleus Definition

As discovered by Ernest Rutherford, atoms are composed of negatively charged electrons orbiting around a very dense region of positive charge. This region of charge is known as the Nucleus of the atom and contains more than 99.9% of the mass of the atom

As the nucleus is positively charged, it produces an electrostatic force that attracts the negatively charged electrons keeping them bound within the atom. Many of the properties of the nucleus determine the properties of the atom as a whole, for example, the size, stability, and radioactivity of an atom are essentially entirely determined by the nucleus.

Nucleus Structure and Function

Unlike electrons, which are fundamental particles, the nucleus is composed of smaller sub-atomic particles known as nucleons. There are two types of nucleons, protons and neutrons; the key properties of these sub-atomic particles are listed in the table below.

As can be seen, protons are the source of the nucleus' positive charge, with each proton having an equal but opposite charge to an electron. Hence, in neutral atoms, the number of protons in the nucleus always equals the number of electrons orbiting the atom. You may be wondering how it is that a nucleus stays stable, given that it consists of positively charged protons packed together. These protons exert electrostatic repulsive forces on each other, so how does the nucleus not fly apart under this force? It turns out there is a second fundamental force involved, known as the Strong Force. This force acts between protons and neutrons, keeping them bound together, this force is the strongest of all the fundamental forces, around \(100\) times stronger than the electromagnetic force pushing the protons apart and so the nuclei can remain densely packed. This explains the role of neutrons within the nucleus, as chargeless particles, they do not affect the electrostatic repulsion between protons; however, they do increase the strength of the Strong Force attraction between the nucleons.

This balance between the electrostatic repulsion between the protons and the strong force attraction between the protons and neutrons determines whether a nucleus is stable. Large nuclei with many protons, known as proton-rich nuclei, become unstable as the electrostatic repulsion starts to outweigh the strong force attraction. This is why heavy elements like Uranium or Plutonium are radioactive as they emit radiation in an attempt to stabilize themselves.

Atoms are characterized by the number of protons and neutrons in their nucleus in the following way. The atomic number \(Z\) determines which element an atom belongs to and is equal to the number of protons within the nucleus. The mass number \(A\) gives the total number of nucleons, that is protons and neutrons, within the nucleus. This is usually denoted as \[\ce{_Z^AX}\]

where \(X\) is the symbol for the element. For example, carbon-12 contains \(6\) protons and \(6\) neutrons and is denoted \[\ce{_6^12C}\]

Atoms of the same element with the same atomic number but a different mass number are known as isotopes.

As neutrons are chargeless, different isotopes have equivalent chemical properties however different isotopes can have different radioactive properties depending on whether the ratio of protons and neutrons is stable or not. For example, Carbon-12 is completely stable however Carbon-14 \(\ce{_6^14C}\) is a radioactive nuclei.

Size and Structure of Atomic Nuclei

An atomic nucleus is an extremely dense object, tens of thousands of times smaller than the entire atom itself. Again the size of the nucleus is determined by the number of protons and neutrons within it, not only because the protons and neutrons have a non-zero volume, but also because of the balance between the electrostatic repulsion and strong force attraction. Unstable proton-rich nuclei will be particularly large as the electrostatic repulsion forces the protons apart. This can lead to some unusual shapes for large nuclei similar to that of an American football.

The radius of a nucleus can vary from around \(1.6\times10^{-15}\,\mathrm{m}\) for the single proton Hydrogen nucleus to around \(15\times10^{-15}\,\mathrm{m}\) for large nuclei such as Uranium. In comparison, the radius of an entire atom, including the orbiting electrons, is usually on the scale of \(1\times10^{-10}\,\mathrm{m}\). Given that the electrons themselves are too small to have any meaningful radius, this goes to show that the atom is almost entirely composed of empty space. However, as we shall see when we look at the electron cloud, quantum mechanics shows us that really the electrons are 'smeared' across the empty space, and we can't really think of them as being particles with a fixed size.

Atomic Structure and Electron Configuration

As we now have an idea about the structure of the nucleus and its role within the atom, let's consider the electrons within an atom. Chemical bonds between atoms are due to the transfer or sharing of electrons, so electrons and their structure are entirely responsible for the chemical properties of an atom. The electronic structure of an atom can tell us how readily it forms bonds with other atoms and what type of bonds it will form, and so a key problem in physical chemistry is to understand the electron configuration of atoms more deeply.

The core concept at the heart of electronic structure is that of an energy level, or electron 'shell'. Quantum mechanics states that the angular momentum and energy of electrons must take fixed quantized values, rather than taking a continuum of values as allowed by Classical physics. This means that electrons must always exist within specific 'shells' which determine the electron's energy. When an atom is excited by an energy source, such as electromagnetic radiation, the electrons can 'jump' up an energy level, but only if the incoming energy is equal to the energy difference between the two levels. Similarly, excited electrons can lose some of this energy by dropping down an energy level and emitting a photon with energy equal to the difference in energies between the two levels. This phenomenon of electron transitions gives rise to the discrete emission and absorption spectra of many atoms, where the light emitted or absorbed by atoms only come in fixed discrete wavelengths.

What's more, only a fixed number of electrons can exist within each shell, determined by the number of allowed values of angular momentum at that energy. This leads to a set structure for the possible electron states within an atom, and these electron states are filled up as the number of electrons in an atom increases. Whilst the full explanation of this structure is beyond the scope of this article, let's take a look at some of the key points. This structure is described using four 'Quantum Numbers' characterizing the state of an electron:

Thanks to the Pauli exclusion principle, which states that no two electrons can have the same quantum numbers, we can use these quantum numbers to determine the maximum number of electrons in an energy level.

It's this configuration of possible electron states which gives rise to the structure of the periodic table. In particular, the number of electrons in an atom's outer shell indicates how reactive an element will be. Elements that have 'unfilled' outer energy shells tend to be the most reactive, for example, the highly reactive Group 1 elements such as Lithium and Potassium, all have a single electron in their outer \(s\) orbitals and so are very reactive.

Atomic Structure: Electron Cloud

As touched on in the previous section, whilst it is often easiest to think about electrons as little 'planets' orbiting around the nucleus, with a well-defined orbit the laws of quantum mechanics show us that things are a bit more complex than that. Heisenberg's Uncertainty Principle, one of the fundamental laws of quantum mechanics, proves that we cannot know the precise location of an electron around a nucleus.