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History and development of atomic theory

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Intros
Lessons
  1. History of the atom
  2. i) Ancient Greece – properties of matter.
    ii) John Dalton: elements and atoms.
  3. JJ Thomson: electrons and the plum-pudding model.
  4. Ernest Rutherford: the nucleus and the Geiger-Marsden (gold foil) experiment.
  5. Niels Bohr: The beginning of quantum theory.
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Examples
Lessons
  1. Recall the scientists involved in the development of atomic theory and their contributions.
    Which scientist's theory was the first to use empirical evidence in forming their ideas?
    a) JJ Thomson
    b) Ernest Rutherford
    c) John Dalton
    d) Niels Bohr
    1. Recall the scientists involved in the development of atomic theory and their contributions.
      Which scientist's theory was the first to use quantum theory in forming their ideas?
      a) Ernest Rutherford
      b) Niels Bohr
      c) JJ Thomson
      d) John Dalton
      1. Recall the scientists involved in the development of atomic theory and their contributions.
        Which scientist's theory was the first to recognise sub-atomic particles?
        a) Niels Bohr
        b) JJ Thomson
        c) John Dalton
        d) Ernest Rutherford
        Topic Notes
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        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 observation, hypothesis and experiment.
        Notes:

        • Today, the words ‘atom’ and ‘element’ are learned very early in chemistry and are well understood. We have a detailed periodic table showing all the elements and a detailed model of the atom and its features.
          But the idea of 'atoms' is thousands of years old. The word comes from the ancient Greek “atomos” meaning indivisible. However, without microscopes and other technology, in ancient Greece the best way of describing matter was by the properties that could be felt by human senses, such as hot or cold and dry or wet.
          These properties were held by the classical elements:
          • Fire was hot and dry while water was cold and wet.
          • Earth was cold and dry while air is hot and wet.
          • There was also “ether”, the substance that fills empty space.
          Many cultures around the world had somewhat similar ideas – the idea that ultimately, all the complex matter in the universe is made up of much smaller, simpler substances or energies that interact with one another.

          Because the technology to measure and see atoms did not exist, not many serious scientific experiments could be done to investigate them. Little progress was made in atomic theory until the 19th century, when a number of experiments were done and their results could be explained by atomic theory.

        • In the early 19th Century, John Dalton proposed his atomic theory: matter came in a variety of elements, and all the atoms of a given element were identical in mass and their other properties.
          These atoms couldn't be destroyed or created, only rearranged and combined in different ways. This became the conservation of mass, which is part of our current understanding of a chemical reaction.
          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 a positively charged cloud to balance out the electrons and give the overall neutral atom.

        • Ernest Rutherford’s gold foil experiment was a massive breakthrough which tested Thomson’s ‘plum pudding’ model. Positively-charged alpha particles were fired at high-speed at a thin gold foil sheet and the way they deflected was recorded.
          • Rutherford’s experiment is a classic example of how the scientific method works in an observation-hypothesis-experiment cycle:
            • Scientists take an observation from an existing topic: according to the plum pudding model, atoms are neutral overall, electrons are small negatively charged particles inside the atom so the rest of the atom must be sparse, dispersed cloud of positive charge.
            • Scientists create a hypothesis or prediction to test the observation. They predict a result that almost no deflection of alpha particles will take place, because if the observation was true then the fast-moving charged alpha particles will not be deflected by the sparse, ‘cloudy’ positive charge of the gold atoms.
            • Scientists design an experiment which applies the hypothesis, where a measurable result will tell you if it is true or not. The hypothesis is about deflection of particles, so the scientists measure deflection of particles. If the hypothesis is true, there would be almost no deflection. If it is not true, there will be significant deflection

          Instead of virtually no deflection in all the alpha particles, while most particles passed through unaffected, some had huge deflection angles. Some particles even scattered back towards the source.
          • To add to the above, the observation (made from the plum pudding model) did not make sense anymore. How could the densely charged high velocity alpha particles get knocked back and away by the sparse charge of the gold atoms? The emprical evidence did not back up the hypothesis made before the test.
            When this happens, the earlier observation is incorrect. In its place, we have the new observation, the results of this experiment. Scientists then need to develop a new theory or revise the current one so that it accounts for the new observation/evidence.

          From this failed hypothesis, Rutherford developed his own atomic theory. What did some of the alpha particles collide with that caused such a huge deflection in their path? Rutherford explained that the atom’s positive charge is entirely concentrated in a tiny core of the atom called the nucleus. This is also where most of the atomic mass is found.
          • Positive alpha particles colliding with a positive nucleus would cause strong charge repulsion and radically deflect the particles from their path. As the nucleus’s size is a tiny fraction of the whole atom, only a small fraction of particles would deflect like this. This is exactly what the experimental evidence showed – the revised theory now correctly explains the new experimental evidence. We are now back at a better observation that is explained by a better theory.
            This is how the scientific method, using an observation-hypothesis-experiment loop, self-corrects and improves understanding. See our lesson CAP.1.1: Using the scientific method for more on this.

          The electrons, relatively tiny, form 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 Rutherford’s planetary model’s weaknesses. If it 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 atomic emission spectra (AES). There was also important new evidence in quantum physics at the time, which Bohr applied to his model.
          The Bohr Model showed electrons being held in discrete shells or energy levels. In the same way, electrons moved up an energy level if a quantized amount of energy was absorbed, and emitted the same quantized amount to move back down to its ordinary (ground) state. This was the first time quantum physics had been used to explain atomic structure. The Bohr model is now obsolete, but electron energy levels in the Bohr model set a foundation for the current understanding of atoms and electron structure.