Why did rutherford model of the atom require that electrons be in motion whereas Thomson's model did not.
The main reason why Rutherford's model of the atom required that electrons be in motion, while Thomson's model did not, is because of their different understanding of the structure of the atom.
Thomson's model, also known as the "plum pudding model," proposed that the atom was made up of a positively charged "pudding" with negatively charged electrons randomly dispersed within it. According to this model, there was no need for the electrons to be in motion because they were assumed to be uniformly distributed throughout the atom. However, this model couldn't explain certain phenomena observed in experiments, such as the scattering of alpha particles.
On the other hand, Rutherford's model, also known as the "nuclear model," proposed that the atom consisted of a small, dense, positively charged nucleus at the center, with electrons orbiting around it. Rutherford's discovery of the existence of the nucleus, through his famous gold foil experiment, led him to conclude that the positive charge and most of the mass of the atom was concentrated in this small region.
To explain the scattering of alpha particles in his experiment, Rutherford hypothesized that the electrons must be moving in orbit around the nucleus. This motion allowed for the alpha particles to interact with the electrons and be deflected. If the electrons were stationary, as per Thomson's model, they would not be able to interact with the alpha particles and cause any scattering.
So, Rutherford's model required electrons to be in motion because it explained the deflection of alpha particles in his experiment, while Thomson's model did not require motion since it assumed that electrons were evenly distributed throughout the atom.