In this lesson, we will learn:
- The key scientists involved in developing atomic theory.
- The steps in the breakthroughs to lead to our modern understanding of the atom.
- How a coherent scientific theory is developed by experiment and observation.
- The idea of the macroscopic world being composed of extremely small indivisible parts or pieces has been present at least since ancient Greek times. The word atom comes from ancient Greek "atomos" – indivisible. Here, the closest attempts to describe matter were by properties such as hot/cold and dry/wet, possessed by the elements fire, water, wind and earth.
- Because there was no technology that allowed people to measure and see such small objects like atoms, no relevant scientific experiments could be performed on the issue. Therefore the issue started largely a philosophical one, not a scientific one. Little progress was made until the 19th Century, when a number of experimental results had emerged that were explainable by atomic theory.
- In the early 19th Century, John Dalton proposed his atomic theory that matter came in a variety of elements, and all the atoms of one element were identical in mass and their other properties. These atoms cannot be destroyed or created in chemical reactions, only rearranged and combined in various ways – a law that still underpins chemical reactions. Dalton also made major contributions to our knowledge of chemical compounds and formulae, measuring the relative masses of elements which he found reacted together to make new chemical substances.
- JJ Thomson is credited with discovering the electron, as a small electrically charged part of an atom. He took the correct idea that atoms are neutral overall and devised the plum-pudding model: electrons were negatively charged 'plum' chunks sitting dispersed through the rest of the atom – the 'pudding', which must be positively charged to balance out the electrons.
- Ernest Rutherford's famous gold foil experiment was a massive breakthrough which tested Thomson's 'Plum Pudding' model. Positively-charged ions were fired high-speed at a sheet of gold foil and the way they deflected was recorded. Expecting that nearly no deflection would happen, Rutherford observed huge ion deflections: some ions scattered backward and other large deflections. If the atom was just a light, spread out 'pudding' of positive charge, how could the positive ions get knocked back and away like that? The observed results of the experiment (the empirical evidence) did not back up the predictions made before the test. The only way to explain the huge backwards deflections was that the atom's positive charge (that caused deflection because of positive-positive repulsion) is entirely concentrated in a tiny core in the middle of the atom, and so was the atom's mass – this was called the nucleus. The electrons, relatively tiny, formed a 'cloud' surrounding the nucleus. This is the planetary model of the atom. Later experiments on nuclear mass led him to conclude the existence of neutrons.
- Niels Bohr dealt with the planetary model's weaknesses. If the model was true, the cloud of negative electrons orbiting a nucleus would continuously lose energy and spiral into the nucleus. It also didn't explain why atoms released light of specific energy when heated (known as discrete emission spectra). The Bohr Model showed electrons being held in discrete shells or energy levels. In the same way, electrons moved between energy levels if a specific or quantized amount of energy was absorbed or emitted. This was the first time quantum physics had been used to explain atomic structure. It's now obsolete, but electron energy levels in the Bohr model set a foundation for the current understanding of atoms and electron structure.