GRAVITATIONAL PUSH AND PULL

In an atom, why doesn’t the negatively charged electron collapse into the positively charged nucleus?

Is this in any way similar to the reason why large systems like stars and planets do not collapse into each other under the pull of gravity?

WHEN Ernest Rutherford, the New Zealand-born founder of nuclear physics, first discovered the atomic nucleus he did suggest that electrons did not fall toward the nucleus of the atom. This was because, he said, that the attractive forces of the nucleus were being balanced by the orbital velocity of the electron, in much the same way as a planet orbits a star.

However, the Danish physicist Niels Bohr modified this theory after Albert Einstein and Max Planck found that energy could only exist in certain discrete amounts, or quanta. This meant that electrons could be seen to have both wave and particle properties, and required that the circumference of the orbit of an electron could not be zero. This means, of course, it could never reach the nucleus.

Science has since adopted the model by the Austrian and Nobel Prize-winning theoretical physicist Erwin Schrödinger. Instead of orbiting the nucleus like planets, his model has electrons occupying ‘clouds’ where it is statistically probable that they will exist, although it has to be appreciated we may never determine an electron’s position and velocity at the same time.

Niels Bohr’s insight in 1913 is worth further explanation. The atom was known to have a small heavy nucleus, and the much lighter electrons were thought to orbit it like planets around the sun. As long as a planet does not lose energy, it can continue its orbit indefinitely.

According to the laws of electromagnetism, charged particles moving in a circle ought to radiate energy as waves. Bohr calculated that a hydrogen atom should collapse with a flash of light in a matter of femtoseconds. Because this does not happen, he proposed what has become known as the ‘old’ quantum mechanics. It asserted that the electron’s angular momentum had to be a multiple of Planck’s constant.

The rule meant that electrons could only occupy particular orbits, and there was a minimum size of orbit. Using this, Bohr was able to predict the entire spectrum of excited states of hydrogen, which, of right, was quite an astounding achievement.

But Bohr’s theory was difficult to apply to more complex atoms and was superseded by Erwin Schrödinger’s wave mechanics in 1927. This became the start of modern quantum theory.

Schrödinger’s formulation shows that an electron has a wave character, and a stable atom can be thought of as a box confining the wave. An electron has a wavelength equal to Planck’s constant divided by its momentum, so the faster an electron moves, the shorter its wavelength. To confine the electron near the nucleus the electron must move very quickly.

Conversely, a fast-moving electron can escape the pull of the nucleus. So you can think, too, of the size of an atom as resulting from a compromise between the electrons having enough kinetic energy for their waves to fit in the box, but so much that they can escape.