According to Rutherford's model of the atom, proposed by Ernest Rutherford in 1911, electrons were required to be in motion to prevent them from collapsing into the nucleus.
Rutherford's model depicted the atom as a small, dense, positively charged nucleus at the center, with electrons orbiting around it. This model was based on the famous gold foil experiment, in which Rutherford bombarded gold foil with alpha particles and observed their scattering patterns.
The experiment revealed that most alpha particles passed through the gold foil with minimal deflection, suggesting that atoms were mostly empty space. However, a small fraction of particles experienced large-angle deflections or even bounced back. Rutherford interpreted this result as evidence that the positive charge and most of the atom's mass were concentrated in a tiny, dense nucleus.
If electrons were at rest in this model, their negative charge would attract them towards the positively charged nucleus, due to the electromagnetic force. According to classical electromagnetism, an accelerating charged particle loses energy in the form of radiation and should eventually spiral into the nucleus. This would result in the collapse of the atom.
To avoid this collapse, Rutherford proposed that electrons must be in motion, orbiting the nucleus similar to how planets orbit the sun. This motion creates a balance between the attractive force of the positively charged nucleus and the centrifugal force of the electron's motion. The electrons remain in stable orbits without continuously losing energy and collapsing into the nucleus.
Rutherford's model of the atom was a significant step in understanding atomic structure, but it had limitations. It did not explain the stability of atoms with multiple electrons, the observed spectral lines, or the details of atomic behavior. These shortcomings were addressed by the development of quantum mechanics, which introduced the concept of electron wave functions and described atomic structure in a probabilistic manner.