Материал: Андрианова English Thesaurus of Student-Physicist 2013

2. Match the words in A with the words in B to form word combinations:

3.Fill in the gaps with the missing words from the list: photons, work function, quantum effects, uncertainty principle, x-

rays

1._____ can be even more bizarre than relativistic effects, but don't despair: confusion is normal and expected.

2.Photoelectrons are created by absorption of a single photon, so the energy of that photon must be greater than or equal to the

_____, else no photoelectrons will be produced.

3.We now know that _____ are a part of the electromagnetic spectrum, characterized by frequencies higher than those of ultraviolet radiation and having the ability to penetrate most materials with relative ease.

4._____ have wave and particle characteristics, perhaps all forms of matter have both properties.

5.To understand the physical origin of the _____, consider the experiment introduced by Heisenberg.

UNIT 28

ATOMIC PHYSICS

Read the following texts. Study the thesaurus on fig. 28, answer the questions after the texts

In this Unit we first discuss the Bohr model of hydrogen, which helps us understand many features of that element but fails to explain finer details of atomic structure. Next we examine the hydrogen atom from the

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viewpoint of quantum mechanics and the quantum numbers used to characterize various atomic states. Quantum numbers aren't mere mathematical abstractions: they have physical significance, such as the role they play in the effect of a magnetic field on certain quantum states. The fact that no two electrons in an atom can have the same set of quantum numbers - the Pauli exclusion principle - is extremely important in understanding the properties of complex atoms and the arrangement of elements in the periodic table. Finally, we apply our knowledge of atomic structure to describe the mechanisms involved in the production of x-rays, the operation of a laser, and the behavior of solid-state devices such as diodes and transistors.

Early models of the atom

The model of the atom in the days of Newton was a tiny, hard, indestructible sphere. Although this model was a good basis for the kinetic theory of gases, new models had to be devised when later experiments revealed the electronic nature of atoms. J. J. Thomson (1856-1940) suggested a model of the atom as a volume of positive charge with electrons embedded throughout the volume, much like the seeds in a watermelon.

In 1911 Ernest Rutherford (1871-1937) and his students Hans Geiger and Ernest Marsden performed a critical experiment showing that Thomson's model couldn't be correct. In this experiment, a beam of positively charged alpha particles was projected against a thin metal foil. The results of the experiment were astounding. Most of the alpha particles passed through the foil as if it were empty space, but a few particles deflected from their original direction of travel were scattered through large angles. Some particles were even deflected backwards, reversing their direction of travel. When Geiger informed Rutherford of these results, Rutherford wrote, "It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15inch shell at a piece of tissue paper and it came back and hit you."

Rutherford explained these astounding results by assuming that the positive charge in an atom was concentrated in a region that was small relative to the size of the atom. He called this concentration of positive charge the nucleus of the atom. Alpha particles themselves were later identified as the nuclei of helium atoms.

Atomic spectra

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The hydrogen atom is the simplest atomic system and an especially important one to understand. Much of what we know about the hydrogen atom (which consists of one proton and one electron) can be extended directly to other single-electron ions such as He+ and Li2+. Further, a thorough understanding of the physics underlying the hydrogen atom can then be used to describe more complex atoms and the periodic table of the elements.

If a voltage applied between metal electrodes in the tube is great enough to produce an electric current in the gas, the tube emits light having a color that depends on the gas inside. (This is how a neon sign works.). When the emitted light is analyzed with a spectrometer, discrete bright lines are observed, each having a different wavelength, or color. Such a series of spectral lines is commonly called an emission spectrum. The wavelengths contained in such a spectrum are characteristic of the element emitting the light. Because no two elements emit the same line spectrum, this phenomenon represents a marvelous and reliable technique for identifying elements in a gaseous substance.

In addition to emitting light at specific wavelengths, an element can absorb light at specific wavelengths. The spectral lines corresponding to this process form what is known as an absorption spectrum. The absorption spectrum of an element has many practical applications. For example, the continuous spectrum of radiation emitted by the Sun must pass through the cooler gases of the solar atmosphere before reaching the Earth. Because the Greek word for Sun is helios, the new element was named helium.

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The Bohr theory of hydrogen and its modification

At the beginning of the 20th century, scientists were perplexed by the failure of classical physics to explain the characteristics of spectra. In 1913 Bohr provided an explanation of atomic spectra that includes some features of the currently accepted theory. The equations of his theory is based on the assumption that the electron can exist only in certain allowed orbits determined by the integer n. The orbit with the smallest radius, called the Bohr radius, corresponds to n = 1. The lowest energy state, or ground state, corresponds to n = 1. The minimum energy required to ionize the atom - that is, to completely remove the electron - is called the ionization energy.

Newtonian mechanics cannot be used to describe phenomena that occur at speeds approaching the speed of light. Newtonian mechanics is a special case of relativistic mechanics and applies only when v is much smaller than с. Similarly, quantum mechanics is in agreement with classical physics when the energy differences between quantized levels are very small. This principle, first set forth by Bohr, is called the correspondence principle.

Within a few months following the publication of Bohr's paper, Arnold Sommerfeld (1868-1951) extended the Bohr model to include elliptical orbits. Bohr's concept of quantization of angular momentum led to the principal quantum number n, which determines the energy of the allowed states of hydrogen. Sommerfeld's theory retained n, but also introduced a new quantum number, called the orbital quantum number. Another modification of the Bohr theory arose when it was discovered that the spectral lines of a gas are split into several closely spaced lines when the gas is placed in a strong magnetic field. This is called the Zeeman effect, after its discoverer. In order to explain this observation, a new quantum number called the orbital magnetic quantum number.

Finally, very high resolution spectrometers revealed that spectral lines of gases are in fact two very closely spaced lines even in the absence of an external magnetic field. This splitting was referred to as fine structure. In 1925 Samuel Goudsmit and George Uhlenbeck introduced the idea of an electron spinning about its own axis to explain the origin of fine structure. The results of their work introduced yet another quantum number, called the spin magnetic quantum number.

For more than a decade following Bohr's publication, no one was able to explain why the angular momentum of the electron was restricted to discrete

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