Bohr’s Quantum Theory of the Atom

In the same way we have refined our understanding of light over the years, scientists have embarked on a similar journey with the nature of matter and the atoms that comprise it. In this tile we will embark on a whistle-stop tour of the most influential atomic models which have been developed, and the experimental evidence that inspired them.

Once we have explored the dramatic shifts in how atoms and subatomic particles have been envisaged over time, we will discover what the building blocks of our universe look like when examined under a quantum lens.

In 1803, English chemist John Dalton published his atomic theory which revolutionized the field of chemistry. He hypothesized that matter is made up of tiny, indivisible particles called ‘atoms’ which could form molecules and be rearranged, combined, or separated during chemical reactions. A giant leap forward, but his theory failed to predict the existence of subatomic particles such as protons, neutrons, and electrons.

English Physicist J.J. Thomson outlined his ‘Plum Pudding Model’ of the atom in 1904 a few years after discovering the electron in 1897. Since his work predated the ‘atomic nucleus’, he represented the atom as negatively charged “plums” embedded in a positively charged “pudding”.

His model attempted to explain the two known properties of atoms at the time: 1) the negative charge of electrons and 2) that atoms have zero net charge, meaning the positive and negative components must balance out somehow. His work stands out as the very first attempt to represent the atomic structure of matter. He visualized materials and substances as being made up of countless small spheres, each sphere having a positive charge spread uniformly around its volume and being distributed with electrons. It wasn’t long however until a new ‘nuclear’ model was announced a few years late.

In 1909, Ernest Rutherford oversaw the ‘Geiger-Marsden experiments’. The experiments involved firing a narrow beam of ‘alpha particles’ emitted from a decaying radioactive source at a very thin sheet of gold foil.

If Thomson’s plum pudding model were correct, the alpha particles would have whizzed through the sheet undeflected.

Alpha particles, as used in Rutherford's experiment, are helium nuclei, meaning they consist of two protons and two neutrons. They carry a positive charge because of the protons. If Thomson's 'plum pudding' model was correct — where an atom is a uniform positive soup with negative electrons embedded within — the positively charged alpha particles should have passed through the thin gold foil without deviation.

While most of the alpha particles did pass through the foil, a surprising number were significantly deflected, or even bounced back. Rutherford used these unexpected results to conclude that atoms must be primarily empty space, which allowed most alpha particles to pass straight through.

The large angle deflections were explained by the presence of a concentrated positive charge within the atom, which repelled the also positively charged alpha particles.

The small fraction of alpha particles that bounced straight back, around 1 in 10,000, led Rutherford to further infer that the positive charge and most of the atom's mass are concentrated in a tiny volume at the center of the atom, known as the nucleus. This deviated significantly from the plum pudding model, leading to our modern understanding of atomic structure.

The electrons in an atom are usually arranged such that its overall energy is kept as low as possible i.e., the ‘ground state’. When an atom is supplied with energy which its electrons absorb to move into higher ‘energy levels’, the atom is said to enter an ‘excited state’.

These states aren’t stable, meaning the atom can’t remain in this configuration for long. Upon returning to the ground state, the electrons return to their original energy levels, emitting energy in the form of electromagnetic radiation which we can detect. Different elements emit different frequencies of light which can be plotted as ‘atomic emission spectra’.

Classical physics had no explanation for these spectra, as instead of being ‘continuous’ and containing all frequencies, they were characterized by discrete ‘spectral lines’. This indicated that the energy released by electrons in excited atoms was quantized, providing further proof of the quantization of light and leading to atomic models based on quantum theory.

Once the existence of a positively charged nucleus was established by Rutherford’s experiments, he put forward a nuclear model which eclipsed Thomson’s plum pudding model as it was far better able to explain observations.

The atom now consisted of a tiny, dense, positively charged ‘nucleus’ surrounded by a cloud of – comparatively much lighter – negatively charged electrons which orbited the nucleus like planets revolving around the sun. Nearly all the mass of an atom was concentrated in its nucleus!

Unfortunately, the model wasn’t without its issues. The most prominent flaw was the fact that this planetary model of the atom failed to explain why the atoms of a given element produce discrete – and not continuous – ‘line spectra’ as there was not yet a concept of quantized electron energy levels. His model also couldn’t explain why the cloud of negatively charged electrons surrounding the nucleus remained in orbit instead of tumbling into the positively charged nucleus that they are attracted to.

Bohr modified Rutherford’s model by stating that the electrical force between the nucleus and orbiting electrons was mathematically identical to the gravitational force responsible for keeping the planets in orbit around the Sun, explaining why electrons don’t simply “fly into” the nucleus.

He also used quantum thinking to posit that only certain electron orbits are permitted in an atom. Since electromagnetic energy is absorbed or emitted if the electrons within an atom move from one orbit – or energy level – to another, Bohr reasoned that the discrete lines seen in atomic emission spectra were due to electrons having specific, quantized energy levels to move between.

Although Bohr’s model could only make poor predictions of the atomic emission spectra of atoms heavier than hydrogen and would need later refinement to better explain some of the more sophisticated phenomena seen in atoms, its use of quantization set the stage for a wealth of exciting developments in physics.

The major issue with Bohr’s model was that it treated electrons as particles moving in precisely defined orbits. Since light demonstrated both particle-like and wave-like properties, who’s to say particles too don’t possess a similar “fuzziness”? What if we could be justified in some situations to treat particles as ‘matter waves’?

Instead of electrons existing in clean orbits, it turns out that each electron is far better described as a ‘cloud’ of possible locations in space. So how do chemists and physicists approximate the location of an electron, if they can no longer tell exactly where it is or how it moves?

The ‘wave function’ which we discuss in the fourth tile can be solved for electrons to give rise to distinctively shaped – depending on the energy level the electron is in – regions within an atom which enclose where the electron is likely to be around ninety percent of the time! These are known as ‘atomic orbitals’.