Further,he proposed that atoms combine in simple whole-number ratios to form compounds; and during chemical reactions, atoms can undergo rearrangement, separation, or combination.
â— The atom was found to be divisible with the groundbreaking discovery of a sub-atomic particle, the electron, in 1897. However, the atom remains the smallest particle that takes part in a chemical reaction.
â— The discovery of isotopes in 1913 proved that the atoms of the same element may not be identical.
â— The discovery of isobars in 1918 proved that atoms of different elements may be the same.
â— In the case of complex compounds, atoms do not combine in simple whole-number ratios.
Despite these limitations, Daltonâ€™s model proved to be a significant stepping stone towards understanding atomic structure.
J.J.Thomson, who discovered the electron in 1897, proposed the Plum Pudding Model of the atom based on a neutral atom composed of negative and positive charges, as illustrated.
As seen in the figure, Thomson believed that electrons were like â€˜plumsâ€™ embedded as point negative charges in a positively charged â€˜pudding,â€™ so it was called the â€˜plum puddingâ€™ model.
â— The model does not explain the emission of the electron spectrum and the light spectrum from the atom.
â— It did not explain the characteristic deflection of Î±-particles.
Rutherfordâ€™s Model (1909)
Rutherford carried out a series of experiments in which he bombarded Î±-particles on a gold foil and measured their scattering pattern by a fluorescent screen placed around it.
â—Most of the Î±-rays passed through the gold foil undeflected or with minimum deflection. This suggested that an atom possesses an abundance of space.
â—Very few rays were deflected with an angle greater than 90Â°, and even fewer reflected the source, which suggested that the positively charged Î±-rays collided with the positively charged central core.
Thus, Rutherford described the atom as having a positively charged core called a nucleus which is surrounded by negatively charged particles called electrons.
â—The model does not explain the distribution of electrons in spatial orbits.
â— The model fails to explain the stability of the atom.
Bohrâ€™s Model (1913)
To overcome the limitations of Rutherfordâ€™s model, Niels Bohr incorporated quantum theory.
â—He introduced the concept of stationary orbits around the nucleus, along which the electrons revolve.
â—Each orbit is of a certain energy level and the electrons in each level revolve stably, without emitting radiation. The orbit closest to the nucleus is of the lowest energy and the outermost orbit has the highest energy.
â—An electron may absorb incident radiation and get excited to a higher energy level and further emit the radiation to fall back to its original orbit. Thus, electrons jump from one allowed orbit to another by gaining or losing energy. This energy difference between the two stationary orbits is given by:
âˆ†E = E2 â€“ E1 = hðœˆ
Where Planckâ€™s constant is represented by â€œhâ€ and the frequency of radiation is represented by â€œðœˆâ€.
â— The attraction between the nucleus and electrons gives rise to centripetal force. This force is equal and opposite to the centrifugal force of electrons moving in orbits.
â— The angular momentum of an electron moving in orbits is the integral multiple of nh/2ð¿, where n is quantum number.
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