Planck imagined that the particles in the bodies oscillate - similar to spiral springs - and absorb and release radiation. The oscillation frequency should correspond to the energy content of the particles. Physical consideration on the one hand and mathematical testing on the other led him finally to a formula that reproduced the data surprisingly well. Everything went together beautifully - except for a triviality: Planck had to insert into the formula a constant whose meaning was completely unclear to him. This quantum of action (h = 6, 626o10-34 JouleoSeconds) indicated that a radiation particle can not absorb any amount of energy, but only individual "packets", so-called quanta. "It was a purely formal assumption, and I did not think much of it, just that I wanted to produce a positive result in whatever circumstances, whatever it might be, " Planck later said. It was an act of despair. When Planck had finished his lecture on December 14, 1900, at the Physics Institute on the Berlin Reichstag bank, the exquisite public must have looked quite perplexed. What should one think about it? Nature made no jumps. Planck himself tried for several years to incorporate his quantum effect in classical physics. In vain. The true nature of this tiny number remained unclear to him. Its fundamental importance was discovered only in the twenties. The day of his talk today is considered the beginning of this physical revolution. The Planck effect quantum combines the energy of a body or gas with the wavelength of the emitted light: the higher the energy or temperature, the shorter the wavelength. This explains why a piece of iron shines with increasing temperature first red (large wavelength) and finally blue (short wavelength). Quantum theory does not explain why the Planck constant has the value as measured in the laboratory today. He must be accepted as natural.
But what if the Planck constant was larger or smaller? The impact on nature would be huge. Suppose the constant is half or twice as large. Then the sun, assuming it still had a surface temperature of 5, 800 degrees, would no longer appear yellow. Rather, it would then release most of the radiation at shorter or longer wavelengths than today. It would appear violet or infrared. In the second case, we would not be able to see our daylight, at least in the present condition of our eyes. Probably, however, the evolution of living beings would have been different: plants today may not be able to use photosynthesis in the usual way with red light and would therefore not be green. The eyes of the living things would have developed differently and would be able to see UV or infrared radiation.
Crystalline spheres in the atom
Although Planck had discovered the central magnitude of quantum mechanics, the connection with the atomic world was only slowly becoming clear. Ernest Rutherford had shot helium nuclei on wafer-thin gold foil and found that most of them were flying unhindered. Only a few were distracted, occasionally bounced back like ping-pong balls. Rutherford thought long and hard about this result until he had the explanation: He enthusiastically shouted to his assistant Hans Geiger, "I now know what an atom looks like!" He thought it was largely empty. At the center there should be a small nucleus surrounded by electrons at long intervals. Although this model from 1910 explained Rutherford's experiment, it posed a crucial problem: According to the well-known laws of electrodynamics, an electrically charged particle emits radiation in a circular path. As a result, it loses energy and approaches the core on a spiral path. It quickly became clear that if Rutherford were right, all the electrons would have to fall into their cores in fractions of a second, according to the laws of classical physics. There are no stable atoms. display
Three years later, one of Rutherford's students radically broke the riddle. Niels Bohr claimed that the electrons would not run on arbitrary orbits around the nucleus, like planets around the sun. Rather, they are assigned paths that they can only leave under certain circumstances: When they pick up a particle of light, they hop onto a path further away from the core, and if they give one, jump back again. As long as the electron remains in its orbit, it is in a stationary, unchanging condition.
Bohr's theory was grotesque. For a long time he had to argue with Rutherford about the publication: "It seems to me you assume that the electron knows from the start where it will stop, " Bohr's teacher countered provocatively. In fact, Bohr seemed to reintroduce the spheres of astronomy into atomic physics. Nevertheless, it can not be denied that his theory could explain the spectrum of luminescent hydrogen gas. And as Einstein heard of Bohr's work, he is said to have said, "This is one of the greatest discoveries." But it was only the beginning.
It has long been known that in some experiments electrons do not appear as small spheres, but clearly show the properties of waves. In 1924, the French prince Louis Victor de Broglie concluded that electrons lead a double life. He suggested that the electron should be thought of as a standing wave surrounding the atomic nucleus. This wave must be such that the orbit circumference always corresponds to an integer multiple of the wavelength.
The Austrian Erwin Schr dinger took up this idea and worked it out mathematically. In 1926, he achieved his breakthrough: in the Schr dinger equation named after him, the electron is a spatially extended wave surrounding the atomic nucleus, which oscillates similarly to a balloon filled with water t. Only certain modes of vibration are possible, with each form corresponding to a particular energy of the electron. In the transition from one form to another, the electron picks up or emits a quantum of light. That explains why an electron does not penetrate into the atomic nucleus. Of course, it did not explain that every now and then electrons also appear as particles.
In fact, the particle somehow hides in the wave. Their intensity indicates the probability of their being in a certain place. This theory contradicts classical physical law, according to which a particle is either in a certain place or not. Exactly what distinguishes quantum mechanics from classical physics: In the realm of atoms only probability statements are possible. Blurred tracks Werner Heisenberg came to the most profound aspect of this atomic blurriness in 1927. At the age of 25 he had recently become Germany's youngest full professor at the University of Leipzig. He had developed a theory competing with Schr dinger to describe the atoms and quantum jumps of the electrons. He stumbled upon the peculiar fact that the location and momentum of an electron can not be measured sharply at the same time. If you determine the exact location, the impulse measurement becomes inaccurate - and vice versa. Planck's quantum of effect indicates the magnitude of this uncertainty.
The Heisenberg uncertainty relation says nothing about the skill of experimental physicists. It describes a fundamental peculiarity of the microworld and has its cause in the fact that particles also appear like "smeared" waves. "If there ever was an experiment that allowed momentum and location to be determined simultaneously and accurately, quantum mechanics would necessarily be wrong, " Heisenberg said. To date, no one has succeeded in such an experiment.
This indeterminacy has a huge impact on nature. It also applies to the couple of time and energy. The shorter the process of measuring the energy of an electron, the less accurate the value becomes. This has a major consequence: in classical physics, no particle can get energy out of nothing. In quantum mechanics, it is at least possible for an electron to borrow energy and return it within the time frame set by the uncertainty relation: the shorter the considered period, the greater the energy credit. With this additional energy particles are able to skip energy barriers that are actually too high for them - for example in the atomic nucleus.
Inside the nucleus, the positively charged protons are bound together on the one hand by the nuclear force, on the other hand, they repel each other because of the electric force and its charge of the same name. In general, the nuclear power predominates. You can imagine the protons rolling around in a pot with too little energy to roll out of. Now, in the core, two protons and two neutrons can coincide to form a particularly stable helium nucleus, which can escape from the pot. This particle has an energy of four million electron volts outside the nucleus. The energy hurdle at its core is about 30 million electron volts high. The missing amount has borrowed the particle in the short term. Physicists talk about tunneling because it looks like the particle has dug a tunnel through the energy barrier.
This also happens with the radioactive decay, which would not exist without the Heisenberg blur. The reverse process, the fusion of two atomic nuclei, is also possible only through the tunnel effect. In the interior of the Sun, for example, the temperature and thus the energy of the hydrogen nuclei is far too low for these positively charged particles to overcome the electrical repulsion force and unite. This allows only the tunnel effect. Without quantum mechanical blurring, there would be no fire inside the sun.
Meanwhile, the tunneling effect is used in technology, for example in the scanning tunneling microscope, with which individual atoms can be examined. An extremely fine metal tip is guided over a surface as close as possible to it. Normally the space between the tip and the surface forms an insurmountable barrier for electrons. The tunnel effect, however, allows again and again particles the seemingly impossible jump. The closer the needle is to the surface, the greater the tunnel current. From its strength, the surface shape can be determined up to atomic resolution.
Technically, the tunnel effect is also used in so-called Josephson junctions, in which two superconductors are separated by an insulator. Nevertheless, electrons that migrate in pairs in superconductors succeed in tunneling through this obstacle. This process can be influenced by magnetic fields. This enables the construction of SQUIDS (Superconducting Quantum Interference Devices), which measure even the weakest magnetic fields, such as brain waves.
Place Cards in the Atom Just two years before Heisenberg's groundbreaking discovery, another intriguing peculiarity of nature had become clear. Bohr had wondered: An atom is usually in the lowest energy state. Should not all electrons have the lowest energy level in this state? Obviously this is not the case, because if all the electrons in the atoms of all elements were in the same state, then all the atoms would have nearly the same appearance. Consequently, there would not be a wealth of chemical compounds and no life.
The Austrian physicist Wolfgang Pauli came to the idea in the beginning of 1925 that nature must have established an exclusion principle. It should dictate that all electrons in an atom occupy different states of energy. In Bohr's atomic model, this means that the shells were filled up with increasing number of electrons from the inside to the outside. Later Pauli transferred his principle to the more abstract quantum mechanical atom model of Schrödinger and Heisenberg.
This suddenly explained the diversity of the elements. Because of the Pauli ban, the atoms in the outermost shell had different numbers of electrons, which are responsible for the chemical properties of an element.
The Pauli principle in solid-state physics has far-reaching consequences. In a crystal, for example, the atoms are so close that their electrons get in the way. As a result, a collective exclusion principle applies to all electrons. That is, no two electrons in the body may be in the same physical state. This has the consequence that the energy levels of the electrons move against each other. They are so close together that they form a broad energy band. In fact, two bands are created. In the low energy band, the electrons are trapped while in the body they can move in the higher energy band and conduct electrical current. The distance between the two bands determines how easily the electrons can jump from the lower band to the higher conduction band. This is the crucial property that differentiates electrical conductors, insulators, and semiconductors.
This nature-imposed prohibition of uniformity of electrons, which Pauli discovered three quarters of a century ago, can not be explained in more detail. It must be accepted as given. Bohr was thrilled with this "complete madness". Pauli, too, did not know what to say: "We must not want to shake the atoms in the shackles of our prejudices." The exclusionary ban is the key to understanding diversity in nature, and it is the starting point for a technical development that goes with construction of the first transistor began and today has reached a peak in microelectronics.
Thomas Bührke Read also from the September issue of bild der wissenschaft The Ultimate Computer Revolution.