Heisenberg's principle:
Werner Heisenberg was a theoretical physicist, who worked with Niels Bohr in Copenhagen in the 1920's. One day he had a shocking realization about the limits of experimental knowledge: the act of observing alters the reality being observed, at the subatomic level. To measure the properties of a particle (such as an electron), one needs to use a measuring device, usually light or radiation. However, the energy of this radiation affects the particle being observed! If you adjust the light beam to accurately measure position, you need a short-wavelength, high-energy beam. It would tell you position, but its energy would throw off the momentum of the particle. Then, if you adjust the beam to a longer wavelength and lower energy, you could more closely measure momentum, but position would be inaccurate.

This epiphany formed the basis of his uncertainty principle, which stated that the accuracy with which one could know both the position, x, and the momentum (momentum p = mass × velocity) of a particle was limited (where Planck's constant, h = 6.57 × 10 erg-sec) to:

h   /   4     x   ×   p

This principle punctured the centuries-old belief that the universe, and everything in it, operated like clockwork (much like the Copernican model of the solar system, in its day, shook up the belief that the Earth lay at the center of the universe). To predict the workings of the clock, one needed only to measure its qualities and parts at a specific point in time. Classical mechanics (the old style of physics) assumed that the precision with which one could measure quantities had no limit. But Heisenberg stated that since you could never measure more than one property of a particle with great certainty, you could only work with probability and mathematical formulations.

Quantum mechanics had several important ramifications. We shall concentrate upon two of them.

• Energy is quantized into small but finite packet sizes. Beyond a certain point, you cannot divide energy into smaller and smaller chunks.

  Consider the orbital paths and energy levels of electrons within atoms. Electrons cannot occupy an arbitrary energy level, but only a few allowed states at set intervals. Classical mechanics assumed that the electron was free to move throughout the atom, with an arbitrary amount of energy at any time. As an analogy, consider a woman walking up the side of a mountain. She can stop at any point, go back a half-step, jump upwards atop a small prairie dog mound, and stand at any height that she chooses. It takes energy to resist gravity and walk up the mountain, so we can relate her height at any time to an energy state. If the woman were operating in the regime of quantum mechanics, however, her path would be very different. She would find herself on a ladder attached to the side of the mountain, able to stand only on the positions were there were steps on the ladder. If she did not have the energy to advance an entire step on the ladder, she would have to stay where she ways (or drop down a full step level).

• Matter does not behave as we expect it to act, at the small distances comparable to the radii of particles (or the wavelength of high-energy light).

  Our conception of how matter behaves is firmly based upon the behavior of matter as we can observe it, which means that it is heavily biased towards people-sized objects. The classical laws derived by Newton which govern the way that gravity attracts and binds masses together, or the trajectories of objects which have been accelerated and are shooting through space, do not necessarily apply to sub-atomic particles or to light.