From the classical models of the 18th century to the quantum revolution of the 20th century, our understanding of atomic structure has evolved. With its probabilistic nature, the quantum mechanical model has fundamentally reshaped our perception of atoms. This article explores the origins and fundamentals of this groundbreaking model.
Image Credit: Marko Aliaksandr/Shutterstock.com
Historical Development of Atomic Models
The quest to comprehend atomic structure began in earnest in the early 1800s. British scientist John Dalton proposed an early atomic theory, picturing atoms as tiny, solid, indivisible spheres. This was followed in the late 19th century by J.J. Thomson's "plum pudding" model, which depicted the atom as a diffuse cloud of positive charge with electrons embedded throughout like plums in a pudding.
However, a key experimental result in 1909 upended established conceptions of atomic structure. Physicist Ernest Rutherford directed alpha particles at a thin gold foil and detected that a small fraction rebounded backward. This suggested that the positive charge and nearly all the mass of an atom is concentrated in a tiny, dense nucleus, contrary to the plum pudding model.
Rutherford conceived the atom as a compact central nucleus orbited by electrons at a distance, akin to planets revolving around the sun. However, this raised a problem: according to classical electromagnetism, the orbiting electrons should continuously lose energy and fall into the nucleus.
Around this time, the discovery of quantized energy levels and emission spectra pointed to deficiencies in classical physics when applied at atomic scales. To address this, physicist Niels Bohr integrated the newly proposed quantum concepts into Rutherford's model. He postulated that electrons occupied discrete circular orbits at set distances from the nucleus, jumping between fixed energy levels and emitting photons when changing orbits.
While Bohr's model successfully explained patterns in atomic spectra and the stability of matter, issues remained. It relied heavily on classical notions of definite electron orbits that did not fit with emerging quantum theory. The model also could not explain finer details of atomic spectra and failed to account for electron properties like spin and magnetic moment.
As a result, the stage was set for a new model of atomic structure based firmly on the principles of quantum mechanics.
The Schrödinger Equation and Quantum Mechanical Model of Atoms
In contrast to Niels Bohr's conception of electrons as discrete particles orbiting the atomic nucleus in fixed paths, Erwin Schrödinger, inspired by de Broglie's particle-wave duality, postulated that the behavior of electrons could be mathematically described as matter waves. This model, known as the quantum mechanical model, laid the foundation for modern understanding of the atomic structure.
The quantum model describes the electrons probabilistically, using Schrodinger's wave equation to determine the likelihood of finding an electron at any given location.
Here, H represents the Hamiltonian operator, which defines the total energy of a quantum particle, and E is the particle's actual total energy. The wave function ψ at a point in space represents the electron's matter wave amplitude, though it often contains complex values.
However, the square of the magnitude of a wavefunction, |ψ|², is significant as it describes the probability of finding an electron within a specific region, serving as the probability density. Schrödinger identified these regions of space as electron orbitals, which differ from Bohr's circular orbit/planetary model.
In Schrödinger's quantum model, electrons are more accurately described as existing in "clouds" around the nucleus, organized into shells or orbitals at varying distances. The Schrödinger equation solutions (quantum numbers) provide the energetics and spatial distribution of electrons in orbitals.
Similar to Bohr's model, the energy of an electron in an atom is quantized with specific allowable values. The key distinction between Bohr's model and Schrödinger's approach is that Bohr introduced quantization arbitrarily, whereas in Schrödinger's approach, quantization naturally arises from the wave equation.
This quantum model effectively addressed inconsistencies in the classical Rutherford-Bohr theory and offered improved agreement with spectroscopic experiments. The quantum mechanical framework also provided a natural explanation for the intrinsic properties of electrons, including spin and magnetic moment, which arise from the underlying wave-particle duality.
Quantum Numbers and Atomic Orbitals
The quantum mechanical model uses four quantum numbers to define the atomic orbitals and their electron distribution.
- Principal Quantum Number (n): Describes an electron's energy level and average distance from the nucleus. It takes positive integer values (1, 2, 3, etc.), with higher values indicating higher energy levels.
- Angular Momentum Quantum Number (l): Determines the shape of the orbital, whether it is a spherical s-orbital, dumbbell-shaped p-orbital, clover-shaped d-orbital, and so on
- Magnetic Quantum Number (m₁): Indicates the spatial orientation of orbitals and varies from -l to 0 to +l based on the value of l.
- Spin Quantum Number (ms): Describes an electron's spin in a magnetic field as either +1/2 (clockwise) or -1/2 (counterclockwise), with each subshell accommodating two electrons with opposite spin values.
Electron Spin and Pauli Exclusion Principle
Austrian physicist Wolfgang Pauli established a fundamental principle, the Pauli exclusion principle, which provides the final information needed to comprehend the overall behavior of electrons within atoms.
This principle is based on the electron spin quantum number, representing the electron's intrinsic spinning motion. It restricts how electrons can be configured within an atom, stating that no two electrons in the same atom can possess an identical set of all four quantum numbers. Therefore, each orbital can contain a maximum of two electrons with opposite spin quantum numbers +1/2 and -1/2.
Concluding Remarks
Despite its complexities, the quantum mechanical model excelled in explaining observations, notably in the hydrogen spectrum, where quantized orbital energies derived from quantum numbers produced multiple spectral lines. This model extended to multi-electron atoms, departing from classical particle-based views to embrace probabilities, wavefunctions, quantized energies, spin, and quantum numbers as fundamental attributes governing particles' behavior, shaping atomic and molecular structure. This transformation is a remarkable achievement in humanity's quest to understand the atom, serving as the foundation of chemistry and materials science.
New Particles and Physics Beyond the Higgs Boson
References and Further Reading
Anderson, G., MacDonald, J., Richards, A., Sullivan Sauer, S., van Haarlem, J. R., & Wegman, D. (2023). 10.4 Quantum Mechanical Model of The Atom - Enhanced Introductory College Chemistry. https://ecampusontario.pressbooks.pub/enhancedchemistry/chapter/quantum-mechanical-atom/
Angelo State University. (2023). Quantum Numbers, Atomic Orbitals, And Electron Configurations. [Online]. Available at: https://www.angelo.edu/faculty/kboudrea/general/quantum_numbers/Quantum_Numbers.pdf
Zumdahl, S. S., Zumdahl, S. A., & DeCoste, D. J. (2016). Chapter 7- Atomic Structure and Periodicity - Chemistry. Cengage Learning. https://www.cusd80.com/cms/lib/AZ01001175/Centricity/Domain/2155/Zumdahl%20Chapter%207%20Atomic%20Structure%20and%20Periodicity.pdf
Disclaimer: The views expressed here are those of the author expressed in their private capacity and do not necessarily represent the views of AZoM.com Limited T/A AZoNetwork the owner and operator of this website. This disclaimer forms part of the Terms and conditions of use of this website.