The history of quantum mechanics – Part II

So as promised… I’m back with the next discussion. If you remember, we were to discuss the prologue of our quantum drama..I mean the riddles which compelled the physicists and chemists to think of a wider explanation.

Before quantum theory was conceived, physicists were unable to explain some of the most ubiquitous phenomena in our environment. Here are a few examples:
A piece of iron, when heated, becomes first red, then yellow, then white, but no body could explain why. The different colors emitted by a piece of matter come from the irregular heat motion of the electrons in the atoms of iron. Fast motion emits higher frequencies of light than slow motion. The frequency determines the color of the light. The laws of thermodynamics tell us that any form of motion should receive the same amount of energy at a given temperature, an amount that increases when the temperature rises. Thus we expect only an increase of intensity of the emitted light (i.e.it must glow more bright dark) not a change from red to yellow to white. This change represented an unsolved riddle, somewhat like today’s ignorance of the nature of memory. We use memory constantly, but nobody knows precisely what it is.

We can consider more examples of unsolved questions of that period: Why are copper and silver metals and oxygen a gas? Why do metals have properties so different from other solids, such as rocks? Why do these properties persist, even after heating, melting, evaporating, and subsequent cooling to the original temperature? Why do oxygen atoms bind to hydrogen atoms to form water? Why does sodium gas emit yellow light when heated? Why is it that burning 1 kilogram of coal produces approximately 6000 calories? Why is the size of atoms about a hundred millionth of a finger’s breadth? No one could provide any answers to these questions at the turn of the century.

More problems were generated by the discovery of the electron by J.J. Thomson in England and H.A. Lorentz in the Netherlands at the end of the last century and, in particular, by further discoveries by E. Rutherford and H.G. Mosely at the beginning of this century. They found that atoms consist of a heavy, positively charged atomic nucleus surrounded by much lighter electrons. Because the electric force between the positively charged nucleus and the negatively charged electrons is of the same form as the attraction between the sun and the planets, they concluded that atoms must be tiny planetary systems, with the nucleus sun and the electrons as planets. Moreover, it was found that all elements seem to have this planetary structure and differ only in numbers of electrons. For example, hydrogen has one electron; helium, two; iron, twenty-six; and uranium, ninety-two.

It was difficult to understand how elements that are so different -some are gases, some are metals, some are liquids- differ in their atomic structure by only a few electrons. For example, the element neon, which has ten electrons, is a chemically inactive gas: however, the element sodium, which has eleven electrons, is one of the most chemically active metals. An electron increase of 10 percent completely changes the character of the atom! No one could explain this apparent inconsistency between quantity and quality.

Four observations defied all understanding at the turn of the century.

  1. The color of objects at various stages of heating (red, yellow, white).
  2. The very different specific properties of elements whose number of electrons is almost the same.
  3. The fact that atoms do not change their properties in spite of the many collisions and interactions that they suffer in a gas or in an ordinary piece of matter. They quickly resume their original qualities after the perturbation. Their stability and their ability to regenerate is completely at odds with what we would expect from a planetary system. If our solar system were to collide or pass another star at a close distance its orbits and patterns would be completely changed and it would not return to its original form.
  4. The energy content of the atom is quantized. An atom can assume a series of definite energies only and never a value in between. This most surprising fact was found at the beginning of this century and is also completely foreign to a planetary system. There is no reason why the energy of planetary motion cannot change by arbitrarily small amounts, for example, when a meteorite hits a planet. An atom, however, can accept or lose only definite amounts of energy, those that would change its energy from one of the values in the series to another.

It became clear to prequantum physicists that the analogy between an atom and a planetary system breaks down completely when atomic properties and processes are examined in detail. On the other hand, these observations left no doubt that the atom consisted of a positively charged nucleus surrounded by electrons, which ought to form a planetary system, according to the laws of mechanics known at that time. Everything, including our own bodies, consists of atoms. Obviously, the most urgent problem for physics at that time was to resolve these contradictions and to achieve a better comprehension of the structure and behavior of atoms.
So with this ends our prologue. What were the theories which helped us in making a more pronounced understanding will be our next discussion? Precisely to say we will discuss wave-particle duality in the next article. We are wave and particle at the same time…believe me……..

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One thought on “The history of quantum mechanics – Part II

  1. This paper is interesting and knowledge full. I enjoyed reading your paper. Keep doing great work like this. All the best! Thank you!

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