579 



Thus the wave numbers of the yellow doublet of sodium were represented symbolically : 

 <r= \s — 2Pi,2. More than 30 years passed before these arbitrary symbols could be given 

 any atomic interpretation. 



The concept of atomic energy levels was first clearly stated in 1913 by X. Bohr who 

 postulated (1) that stationary atomic states exist, and (2) that the frequency of atomic 

 radiation is proportional to the difference between two atomic energy states, hv= (Ei — 

 Ei), the proportionality factor being Planck's constant, h. By 1919 the accumulation of 

 singlet, doublet, and triplet terms found in arc and spark spectra barely sufficed to suggest 

 two general laws of spectral structures : (1) the alternation law which states that even and 

 odd multiplicities alternate in successive columns of the periodic chart of the atoms, and 

 (2) the displacement law which states that the spectrum of an ionized atom resembles 

 that of the preceding atom but the analogous lines are displaced toward higher frequencies. 

 Term multiplicities of atoms or ions are thus determined solely by the number of electrons 

 in the atoms, whereas the atomic charge controls the position of the spectrum. These two 

 facts suggested that electrons and protons were involved in the exegesis of atomic spectra. 



The more complex spectra rtsisted all attempts at interpretation until 1922 when M". A. 

 Catalan deliberately set out to discover a new or more general principle in spectral struc- 

 ture. He found in the arc spectra of chromium and manganese terms having five or six 

 levels which combined to produce groups of lines that he called multiplets. In a few 

 years thousands of terms were found in atomic and ionic spectra, and contemporaneously 

 the present quantum theory of atomic energy levels was developed. As a result of these 

 developments the arbitrary symbols that empirical spectroscopy devised for the yellow 

 doublet of sodium were replaced by the following : 



Each and every item of this spectroscopic notation now has definite physical meaning in 

 terms of a vector model of the Rutherford-Bohr atom which is assumed to consist of a 

 minute but massive nucleus (composed of protons and neutrons) with one or more elec- 

 trons circulating about it. The normal number of electrons in any atom is equal to the 

 atomic number, Z : identical with the number of protons in its nucleus. 



Spectral lines result from changes in atomic energies defined by the positions of one or 

 more optical electrons in successive shells and by their orbital and axial momenta, each 

 of which is associated with an appropriate quantum number. In general, the first large 

 change in atomic energy occurs when an electron jumps from its normal shell, represented 

 by the principal quantum number n, to another shell. These principal quantum numbers 

 identify the successive shells of the periodic system and serve as coefficients to the spectral 

 term symbols S, P, D, F, etc. If an electron is moved from its lowest value of n to n = oo 

 the atom is ionized, and the voltage necessary to remove this electron is called the ionization 

 potential. This ionization energy is expressed in wave number (cm -1 ) or in electron volts 

 (ev) as in Tables 623 and 624. Increasing atomic energies are exhibited in absorption 

 spectra, decreasing energies in emission spectra. 



After that due to a change in n, the next largest change in atomic energy is usually 

 one associated with orbital angular momentum symbolized by an azimuthal quantum 



number / having integral values 0, 1, 2, 3, corresponding respectively to the empirical 



term symbols S, P, D, P, . Electrons with / = are called j-electrons, those with 



/=1, /(-electrons, etc. These four / values and the first seven n values suffice to describe 

 the normal electron configurations of all possible atoms and ions. When two or more 



optical electrons are present, their individual orbital momenta U. U are added vec- 



torially to form a resultant L which is restricted by quantum theory to integral values 

 ranging in the case of two electrons from h + U to \U — U\. The types of spectral terms 

 resulting from various simple configurations of electrons are shown in Table 621. 



TABLE 621.— L VALUES AND SPECTRAL TERMS RESULTING FROM 



TWO ELECTRONS 



Electrons L 



SS 



Sp 1 



pp 12 



pd 12 3 



dd 12 3 4 



df 12 3 4 5 



ff 12 3 4 5 6 



{continued) 



SMITHSONIAN PHYSICAL TABLES 



