chemistry

modern atomic theory

Getting to the theory


  1. Daltons theory was based on the premise that the atoms of different elements could be distinguished by differences in their weights. He stated his theory in 1803. The theory proposed a number of basic ideas:
  • All matter is composed of atoms
  • Atoms cannot be made or estroyed
  • All atoms of the same element are identical
  • Different elements have different types of atoms
  • Chemical reactions occur when atoms are rearranged
  • Compounds are formed from atoms of the constituent elements.


He rationalised the laws of chemical combination which were in existence. He made a mistake in assuming that the simple compound of two elements must be binary, formed from atoms of each element in a 1:1 ratio. his system of atomic weights was not very accurate

Despite these errors, Dalton's theory provided a logical explanation of concepts, and led the way into new fields of experimentation.


2. JJ Thomas

assumed that the basic body of an atom is a spherical object containing N electrons confined in homogeneous jellylike but relatively massive positive charge distribution whose total charge cancels that of the Nelectrons

3. Ernest Butherford

publishes his atomic theory describing the atom as having a central positive nucleus surrounded by negative orbiting electrons. This model suggested that most of the mass of the atom was contained in the small nucleus, and that the rest of the atom was mostly empty space. Rutherford came to this conclusion following the results of his famous gold foil experiment. This experiment involved the firing of radioactive particles through minutely thin metal foils (notably gold) and detecting them using screens coated with zinc sulfide (a scintillator). Rutherford found that although the vast majority of particles passed straight through the foil approximately 1 in 8000 were deflected leading him to his theory that most of the atom was made up of 'empty space'.

4. Niels Bohr

applies quantum theory to Rutherford's atomic structure by assuming that electrons travel in stationary orbits defined by their angular momentum. This led to the calculation of possible energy levels for these orbits and the postulation that the emission of light occurs when an electron moves into a lower energy orbit.