Electrons – How Our Perception of This Subatomic Entity Has Evolved

Advancements in our understanding of electrons have paved the way for current chemical and material science research. The behaviour of electrons is embedded in our approach to explore all physical modules, ranging from mechanisms in organic chemistry to solid-state devices. But our perception of the electron has not always been the way it is today - this essay will first explore the discovery and early interpretations of the electron, then delve into the Bohr model and Quantum Mechanical Theories that have emerged later. J.J. Thomson discovered the electron in 1897 when investigating the deflection of cathode rays travelling through a discharge tube. In the discharge tube consisting of low-pressure gas, Thomson applied potential difference between electrodes and observed luminescence on the glass when magnetic and electric fields were applied perpendicularly. The cathode rays were composed of electrons, therefore they were deflected away from the negatively charged region of the aluminium plate. This causes the emission of photons from the glass, producing luminescence. Following this discovery, many interpretations of atomic structure arose, including the well-recognised Thomson Plum Pudding Model, Nagaoka’s Saturnian Model and Rutherford’s nuclear model. In particular, Rutherford’s idea of a small dense nucleus acted as the basis for development of Bohr’s atomic model. Niels Bohr proposed quantisation of atomic energy levels, after recognising that the production of atomic line spectra was due to electron transition between discrete energy levels. Both absorption and emission spectrums are formed from electrons in atoms gaining energy (excitation) followed by de-excitation and simultaneous emission of photons. He models electrons to be localised in circular orbits, each represented by a principal quantum number (n). Arnold Sommerfeld then introduced elliptical orbits, the first realisation of sub energy levels, by using azimuthal quantum number (l) to describe the shape of sub-shells. The next step in evolution was Louis de Broglie’s hypothesis of wave-particle duality. In 1924, de Broglie introduced the idea that all moving matter has a wavelength, which can be demonstrated from observing a diffraction pattern after accelerating electrons through a crystalline structure with atomic spacing similar to the electron wavelengths. This experiment conducted by Davisson-Germer in 1927 provides solid evidence supporting the wave-particle duality hypothesis since diffraction is a wave property while it is only possible to accelerate particles with a mass. Shortly after, Erwin Schrödinger derived the famous Schrödinger equation; solving the time-independent Schrödinger equation will yield an eigenvalue (a constant describing the discrete energy of a sub-shell) times the wavefunction, which portrays the shape of the probability distribution of the electron, an orbital. Moreover, Max Born’s Interpretation that the wavefunction squared gives electron probability density reinforces the concept that atomic orbitals are regions of high electron probability density. Each atomic orbital can be described using a unique set of quantum numbers, which are quantising constants derived from solutions of Schrödinger equation. They include the previously defined Principal Quantum Number (n), Azimuthal Quantum Number (l), while Magnetic Quantum Number (ml) describes orbital orientation and Spin Quantum Number (ms) portrays electron spin. Pauli Exclusion Principle states that electrons in atoms must have four distinct quantum numbers, therefore the two electrons in an orbital must always have a different spin quantum number of + 1/2 or - 1/2. The quantum numbers can then be used to construct electronic configurations, which are in terms of orbitals: sharp (s), primary (p), diffuse (d) and fundamental (f). Furthermore, recall that wavefunction squared gives the electron probability density, so they can be articulated in terms of three-dimensional density plots to help visualise the shapes of orbitals. Indeed, our understanding of electrons has evolved countlessly in physical science. It is difficult to evaluate the validity of historic attempts to understand electronic structure – the Bohr model isn’t necessarily wrong but just lacks depth. On the other hand, it is imperative to realise that atomic orbitals are mathematical descriptions instead of concrete shapes. Our perception of electrons will always evolve, enabling us to broaden possibilities in scientific research.