A Level Chemistry Revision: Physical Chemistry – Atomic Structure

The topics you study at A level chemistry will build on the knowledge you gained in GCSE chemistry, and one of the topics you’ll explore more thoroughly is atomic structure. 

This involves learning how the modern atomic model has developed over time, from Dalton’s atomic theory, to Thomson’s, Rutherford’s, and Bohr’s atomic models, and how these scientists paved the way for Schrödinger’s Quantum Mechanical model of the atom.

Atomic Structure

Central to the development of chemistry as a scientific pursuit, rather than a mere philosophical proposition, is the idea of the atom and the knowledge of the atomic structure.

Since ancient times, people have wondered what could be the most fundamental structure of physical reality, particularly that of matter. Today, we know that the atomic structure involves negatively charged electrons orbiting a positively charged nucleus. We also know that an atom’s nucleus contains both protons (positively charged) and neutrons (negatively charged).

While this knowledge is fundamental in our modern age, it took centuries of scientific research and experiments for us to understand how atomic structure works. 

Although the ancient Greek philosopher, Democritus (460 – 370 BCE), was one of the first thinkers to propose the idea of an indivisible and fundamental particle called the atom (from the Greek word atomos, meaning indivisible), the modern day scientific theory of atomic structure only started in the early 1800s.

Dalton’s Atomic Theory

As you can see from the timeline above, British scientist, John Dalton, was the first to propose a comprehensive scientific theory about the atom in 1803. His atomic theory had four main assertions:

  1. All matter is composed of fundamental particles called atoms and these particles are indivisible
  2. Atoms of the same element are identical in terms of mass and chemical properties
  3. Compounds are combinations of two or more types of atoms
  4. Chemical reactions are best described by the rearrangements of different types of atoms in the reactants

Though revolutionary for its time, there were a few issues with Dalton’s atomic theory. The most notable was his assertion that atoms were indivisible. This was disproved in 1897 when J. J. Thomson discovered the first subatomic particle: the electron. 

Thomson’s Model of the Atom

In 1904, Sir Joseph John Thomson, known as J.J. Thomson, proposed an improved model of the atom that was based on his experiments with cathode ray tubes. This became popularly known as the plum pudding model

In 1897, Thomson discovered that the atoms of metals, like tungsten, that he sealed inside a vacuum tube emit negatively-charged particles when an electric current is applied to them. These particles were called electrons.

The discovery of electrons showed that these subatomic particles can traverse a vacuum without any conducting materials, like metal wires or air. Electrons can also easily be diverted using magnets, and, based on this fact, their mass was computed to be almost zero.

Thomson surmised that an atom must be composed of negatively charged particles that are embedded on a uniform sphere of positively charged matter. Since atoms have neutral charges, he deduced that an equal positive charge provides balance.

However, Thomson’s plum pudding model ultimately failed because of its inability to explain several observations, like the stability of an atom, or the distribution of positive and negative particles. This is where Rutherford stepped in.

Rutherford’s Atomic Model

By 1911, after years of research and experiments on radioactive elements, New Zealand-born physicist Ernest Rutherford discovered alpha and beta particles. The latter are either positrons or electrons, while the former are the positively charged nuclei of helium.

This made it clear that the atom could be further divided into subatomic particles that included more than just electrons. As a result of this revelation, the plum pudding model had to be abandoned and a new model of the atom was born.

Instead of the positive charge being evenly distributed, as Thomson theorised, Rutherford discovered that it was concentrated in the centre of the atom, while the negative charge was more spread out. He deduced this from his famous gold foil experiment.

For this experiment:

  • Rutherford bombarded a piece of gold foil with a beam of alpha particles
  • Most particles passed through and hit the detector on the other side of the foil
  • Meanwhile, some were deflected and others bounced back

This experiment implied that the gold foil was mostly hollow at the atomic level and that there were smaller, but bigger solid particles at the centre.

Rutherford proposed that instead of being static, the electrons revolve around the central mass and charge of the atom like planets revolving around the sun.

Bohr’s Model of the Atom

Although Rutherford’s model was a breakthrough, it had some flaws, the main one being the way it described the motions of the electrons. In 1913, Danish physicist Niels Henrik David Bohr improved Rutherford’s model by reasoning that if electrons were moving like planets, their orbits would be very unstable.

Based on classical mechanics and electromagnetic theory, any charged particle that’s moving on a curved path emits electromagnetic radiation. This means that the electrons would inevitably lose energy and spiral down into the centre of the atom. 

As a result of this logic, Bohr proposed a new atomic model. He postulated that electrons move in orbits with fixed space and energy. The energy of the orbits becomes higher as the electrons get further away from the nucleus. The emission of electromagnetic radiation only occurs when electrons jump from lower energy levels to higher energy levels.

Schrödinger’s Quantum Mechanical Model

Both Rutherford and Bohr assumed that electrons were solid particles that moved around the nucleus. Both scientists also assumed that electrons followed fixed orbital paths. However, some behaviours of the electrons cannot be satisfactorily explained if the electrons were solid particles with a definite volume and location at a specific time.

In 1926, Erwin Schrödinger proposed yet another improved model of the atom. Instead of particles having fixed diameters and fixed orbits, Schrödinger treated electrons as waves. He applied the wave function equations and created the Quantum Mechanical Model of the atom. Here, the electron was viewed as a standing wave that has no definite shape or location.

The Quantum Mechanical Model of the atom remedied the many limitations of Bohr’s model. Schrödinger’s model was able to:

  • Reconcile the Heisenberg Uncertainty Principle with the electron behaviours by assuming that it has no fixed diameter and the exact location cannot be determined if the velocity is known
  • Make better spectral predictions for larger atoms, unlike the Bohr model, which only worked with hydrogen atoms
  • Predict the relative intensities of spectral lines emitted by atoms when heated during spectral analysis
  • Explain the Zeeman Effect. It was able to account for the splitting of the spectral line into several components under the influence of magnetic fields

Schrödinger’s Quantum Mechanical model of the atom is the best atomic model we have so far. It can correctly predict and explain many of the behaviours of the atom, including chemical bonding. This model also assigned orbitals with various energy levels for electron clouds. It’s a precise mathematical model that allows scientists to calculate the energy yields of atoms.

Looking for more chemistry A level resources? Check out our revision guides and other articles.

Disclaimer

The blog on chemicals.ie and everything published on it is provided as an information resource only. The blog, its authors and affiliates accept no responsibility for any accident, injury or damage caused in part or directly from following the information provided on this website. We do not recommend using any chemical without first consulting the Material Safety Data Sheet which can be obtained from the manufacturer and following the safety advice and precautions on the product label. If you are in any doubt about health and safety issues please consult the Health & Safety Executive (HSE).