Structure of matter

1.4. DISTRIBUTION OF ORBITAL ELECTRONS
According to the model proposed by Niels Bohr in 1913, the electrons revolve around the nucleus in specific orbits and are prevented from leaving the atom by the centripetal force of attraction between the positively charged nucleus and the negatively charged electron.
On the basis of classical physics, an accelerating or revolving electron must radiate energy. This would result in a continuous decrease of the radius of the orbit with the electron eventually spiraling into the nucleus. However, the data on the emission or absorption of radiation by elements reveal that the change of energy is not continuous but discrete. To explain the observed line spectrum of hydrogen, Bohr theorized that the sharp lines of the spectrum represented electron jumps from one orbit down to another with the emission of light of a particular frequency or a quantum of energy. He proposed two
P.6
Fundamental postulates: (a) electrons can exist only in those orbits for which the angular momentum of the electron is an integral multiple of h/2ПЂ, where h is the Planck’s constant (6.62 Г – 10-34 J-sec); and (b) no energy is gained or lost while the electron remains in any one of the permissible orbits.

FIG. 1.2. Electron orbits for hydrogen, helium, and oxygen.

The arrangement of electrons outside the nucleus is governed by the rules of quantum mechanics and the Pauli exclusion principle (not discussed here). Although the actual configuration of electrons is rather complex and dynamic, one may simplify the concept by assigning electrons to specific orbits. The innermost orbit or shell is called the K shell. The next shells are L, M, N, and O. The maximum number of electrons in an orbit is given by 2n2, where n is the orbit number. For example, a maximum of 2 electrons can exist in the first orbit, 8 in the second, and 18 in the third. Figure 1.2 shows the electron orbits of hydrogen, helium, and oxygen

atoms.
Electron orbits can also be considered as energy levels. The energy in this case is the potential energy of the electrons. With the opposite sign it may also be called the binding energy of the electron.
1.5. ATOMIC ENERGY LEVELS
It is customary to represent the energy levels of the orbital electrons by what is known as the energy level diagram (Fig. 1.3). The binding energies of the electrons in various shells depend on the magnitude of Coulomb force of attraction between the nucleus and the orbital electrons. Thus the binding energies for the higher Z atoms are greater because of the greater nuclear charge. In the case of tungsten (Z = 74), the electrons in the K, L, and M shells have binding energies of about 69,500, 11,000, and 2,500 eV, respectively. The so-called valence electrons, which are responsible for chemical reactions and bonds between atoms as well as the emission of optical radiation spectra, normally occupy the outer shells. If energy is imparted to one of these valence electrons to raise it to a higher energy (higher potential energy but lower binding energy) orbit, this will create a state of atomic instability. The electron will fall back to its normal position with the emission of energy in the form of optical radiation. The energy of the emitted radiation will be equal to the energy difference of the orbits between which the transition took place.
If the transition involved inner orbits, such as K, L, and M shells where the electrons are more tightly bound (because of larger Coulomb forces), the absorption or emission of energy will involve higher energy radiation.


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Structure of matter