compare between Bohr model ,J.J Thomson and rutherford model
Answers
Answer:
Thomson’s Atomic Model
In 1898, J. J. Thomson proposed the first of many atomic models to come. He proposed that an atom is shaped like a sphere with a radius of approximately 10-10m, where the positive charge is uniformly distributed. The electrons are embedded in this sphere so as to give the most stable electrostatic arrangement.
Rutherford’s Atomic Model
The second of the atomic models was the contribution of Ernest Rutherford. To come up with their model, Rutherford and his students – Hans Geiger and Ernest Marsden performed an experiment where they bombarded very thin gold foil with α-particles. Let’s understand this experiment.
α-Particle Scattering Experiment
Experiment
In this experiment, high energy α-particles from a radioactive source were directed at a thin foil (about 100nm thickness) of gold. A circular, fluorescent zinc sulfide screen was present around the thin gold foil. A tiny flash of light was produced at a point on the screen whenever α-particles struck it.
Based on Thomson’s model, the mass of every atom in the gold foil should be evenly spread over the entire atom. Therefore, when α-particles hit the foil, it is expected that they would slow down and change directions only by small angles as they pass through the foil. However, the results from Rutherford’s experiment were unexpected –
Most of the α-particles passed undeflected through the foil.
A small number of α-particles were deflected by small angles.
Very few α-particles (about 1 in 20,000) bounced back.
Thomson’s model versus Rutherford’s model [Source: Wikimedia Commons]
Conclusions of the α-scattering experiment
Based on the above results, Rutherford made the following conclusions about the structure of the atom:
Since most of the α-particles passed through the foil undeflected, most of the space in the atom is empty.
The deflection of a few positively charged α-particles must be due to the enormous repulsive force. This suggests that the positive charge is not uniformly spread throughout the atom as Thomson had proposed. The positive charge has to be concentrated in a very small volume to deflect the positively charged α-particles.
Rutherford’s calculations show that the volume of the nucleus is very small compared to the total volume of the atom and the radius of an atom is about 10-10m, while that of the nucleus is 10-15m.
Let us understand in three different ways :-
J.J Thomson :
Thomson discovered the electron which was
negative. He knew atoms were neutral so he
reasoned that there must be some positive stuff in the atom to balance the charge of the electron. In the absence of any other information he proposed a simple atom of positive stuff with electrons embedded in it. This is the plum pudding model.
Rutherford :
Rutherford, through his alpha particle scattering experiment showed that alpha particles were mostly undeflected by a very thin metal sheet (he used gold because Gold is very thin). A few alpha particles rebounded back towards the source deflected through more than 90 degrees.
He reasoned using ideas about momentum that what the alpha particles were hitting/coming up against must have a large mass compared with the alpha particles otherwise the alpha particles would have knocked the bits of gold forward and the alpha particles would not rebound but possibly slow down or stop.
Rutherford was able to deduce that the charge must be concentrated in very small volumes - as most alpha particles were undeflected. This small volume contained both the mass of the atom and the charge. His model was that the atom had a small central nucleus containg mass and positive charge. The electrons would orbit around the nucleus and hence be the outside of the atom.
In this, the electron is accelerating as it moves in its orbit, the electron should rAdiate energy away as electromagnetic radiation and spiral into the nucleus. They knew atoms were stable so did not do this, it was a known problem.
Bohr's:
Bohr' model was an attempt to explain the line spectra emitted by excited gas molecules ( very hot/ electric discharge through the gas), gases behaves very differently to solids in this respect
Bohr explained why only certain frequencies of light are emitted by a Hydrogen atom. He said, if you assume that the electron can only orbit at certain positions where angular momentum was some multiple of a basic amount (h/2pi/ from memory!) then you end up with several positions where the electrons can orbit and the energy gap between these positions exactly matches the observed energies of photons emitted by the hydrogen atoms. The thinking was simple conservation of energy.
Energy of electron in excited orbit = energy of electron in lower orbit + energy emitted by atom in a photon.
The moving from one orbit to another was called a quantum jump. The electrons could not orbit in places in between Bohr knew that there was no justification for not allowing the electron to orbit anywhere-he pointed out that this assumption worked in it gave the correct values for the H spectrum.
OM!