Schödinger's Journery to Atomic Orbials

Schrödinger's Eureka Moment

It was a cold, rainy day in Vienna when Erwin Schrödinger, an Austrian physicist, found himself deep in thought. The year was 1926, and Schrödinger was about to make a groundbreaking discovery that would change the course of chemistry and our understanding of the atom forever. As the raindrops rhythmically tapped on his window, Schrödinger pondered the mysteries of the atom and its elusive electrons.

The Limitations of Bohr's Model

Before Schrödinger, the prevailing model of the atom, developed by Niels Bohr in 1913, suggested that electrons orbited the nucleus in fixed paths or orbits[1]. This model, however, could not explain certain experimental observations, such as the fine structure of hydrogen spectra and the intensities of spectral lines, leaving scientists puzzled[2].

Schrödinger's Revolutionary Equation and the Birth of Atomic Orbitals

Schrödinger, influenced by the burgeoning field of quantum mechanics, formulated a mathematical equation that portrayed electrons in atoms as exhibiting wave-like properties[3]. This pioneering equation gave rise to the concept of atomic orbitals – probability distributions of electron locations around the nucleus. In contrast to Bohr's fixed orbits, orbitals offered a more precise and detailed comprehension of electron behaviour, as demonstrated in the graphic below.

S and P Orbitals: A Closer Look

In this new model, electrons are found in regions of space called orbitals. The most basic types of orbitals are s and p orbitals. S orbitals have a spherical shape and can hold a maximum of two electrons, while p orbitals are shaped like a dumbbell and can hold up to six electrons[1]. The shape of the orbitals is crucial in understanding the distribution of electrons in atoms and how they form chemical bonds.

Electron Spin and Pairing

Another key aspect of atomic orbitals is the concept of electron spin. Electrons can be in one of two spin states, either "spin up" or "spin down." When two electrons occupy the same orbital, they must have opposite spins. This pairing of electrons with opposite spins in the same orbital is known as the Pauli Exclusion Principle[4].

The Legacy of Schrödinger's Discovery and Future Exploration

The discovery of atomic orbitals came as a result of Schrödinger's wave equation, which describes the wave-like behaviour of particles in quantum mechanics. This equation led to the development of the current atomic structure theory, known as the quantum mechanical model of the atom, which still relies on the concept of orbitals.

However, atomic orbitals have their limitations when it comes to explaining the bonding in molecules in terms of electron distribution. In future blog posts, we will explore hybridisation and molecular orbital theory, which provide a more detailed picture of chemical bonding and electron behaviour in molecules. Stay tuned for these exciting topics as we continue our Chemistry Chronicles!

References:

[1]: Atkins, P., & de Paula, J. (2014). Physical Chemistry: Quanta, Matter, and Change (2nd ed.). Oxford University Press.

[2]: Whitten, K. W., Davis, R. E., Peck, L. M., & Stanley, G. G. (2013). Chemistry (10th ed.). Cengage Learning.

[3]: Schrödinger, E. (1926). Quantisierung als Eigenwertproblem. Annalen der Physik, 384(4), 361-376.

[4]: Pauli, W. (1925). Über das Gesetz der Geschwindigkeitsverteilung im Paramagnetismus auf Grund der Quantentheorie. Zeitschrift für Physik, 31(1), 373-382.