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The Periodic Law

Although the arrangement of chemical elements in alphabetical order is convenient for inventory purposes, it is more reasonable to arrange them in the order of increasing atomic weights. In doing so, we find rather remarkable regularities which have led chemists to a rational classification of the elements. Arranging the elements in order of atomic weights,* we obtain the following sequence: H, He, Li, Be, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, A, K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, As, Se, Br, Kr, Rb, Sr, Y, Zr, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Cd, In, Sn, Sb, Te, I, Xe, Cs, Ba, La, etc. We notice, first of all, that there is a remarkable regularity in the distribution of noble gases, shown in italics, throughout the sequence: there is only 1 element preceding He, 7 elements between He and Ne, another 7 elements between Ne and A, 17 elements between A and Kr, and another 17 elements between Kr and Xe. Finally, there are 31 elements between Xe and Rn, which is the heaviest known noble gas.

*The careful student will notice that K, Ni, and I are out of order, but, as it was found later, the sequence of chemical properties has priority over atomic weights.

The elements immediately following the noble gases, lithium, sodium, potassium, rubidium, and cesium, are physically and chemically very similar to each other. They are all light, silvery-white metals with high chemical activity. If we drop a small piece of any of these elements in water, it will undergo a violent chemical reaction of the type:

Li + H2O → LiOH + H

Na + H2O → NaOH+ H

etc.

liberating hydrogen and forming the corresponding "hydroxide" with water (structural formula, Li—O—H, etc.). The hydrogen liberated in this reaction often becomes ignited and produces a flame which takes on the characteristics color of the vaporized metal (yellow for sodium, red for potassium, etc.). Uniting with hydrogen and oxygen, these elements form "hydrates" and "oxides" of the type LiH (Li—H), Li2O (Li—O—Li), etc., showing that their valency is 1. These elements are commonly known in chemistry as alkali metals.

The second neighbors to the right of the noble gases, beryllium, magnesium, calcium, strontium, barium, and radium, also form a homologous group known as alkali-earth metals. As their name indicates, they are similar to the alkali metals, but, as a rule, they are much harder and less reactive. Reacting with water, they produce compounds of the type Ca(OH)2 (H—O—Ca—O—H), while uniting with hydrogen and oxygen they give rise to compounds such as CaH2(H—Ca—H) and CaO (Ca=O), which indicates that their valency is 2. Similarly, we find that the third group to the right, boron, aluminum, etc., possesses a valency of 3 as demonstrated by such compounds as boron oxide, B2O3(O=B—O—B=O), and aluminum hydroxide, Al(OH)3.

Now if we look at the elements standing to the left of the noble gases, we will find that they are very similar to each other, but as different from metals as they could possibly be. This group comprises fluorine, chlorine, bromine, iodine, and astatine, and they are known as the halogens. They have a strong affinity for both alkali and alkali-earth metals, with which they form such compounds as NaCl (ordinary table salt) and CaBr2, indicating that they possess a single valency. The second neighbors to the left of the noble gases, oxygen, sulfur, etc., are also in some ways similar to each other and possess a valency of 2.

The existence of homologous groups and of a certain periodicity in the chemical properties of elements arranged in the order of increasing atom weights was noticed by several chemists during the nineteenth century, but the most important step of actually arranging the elements into a periodic table was made in 1869 by the Russian chemist, Dmitri Mendeleev (1834-1907). Mendeleev was handicapped in his studies because in his time the list of known chemical elements was rather incomplete and, in particular, the existence of the noble gases was not even suspected. From the sequence given above, Sc, Ga, Ge, Tc, and Rh were missing, making the sequence quite irregular except for the first two periods. Driven by a deep belief that there must be a regular periodicity in the natural sequence of elements, Mendeleev made the bold hypothesis that the deviations from the expected periodicity in his list were due to the failure of contemporary chemistry to have discovered some of the elements existing in nature. Thus, in constructing his table, he left a number of empty spaces to be filled in later by future discoveries. He gave to the "missing elements" names formed by adding the prefixes eka or dvi, meaning "first" and "second" in Sanskrit, to the names of neighboring homologous elements. In certain instances, he also reversed the atomic-weight order of elements in order to comply with the demands of the regular periodicity of their chemical properties. Using his table, shaky as it was, he was able to predict the physical and chemical properties of six "missing elements" on the basis of the known properties of their alleged neighbors. He called these elements eka-boron, eka-aluminum, eka-silicon, eka-manganese, dvi-manganese, and eka-tantalum. His predictions turned out to be in excellent agreement with the actually observed properties of the "missing elements" when they were finally found and named: scandium, gallium, germanium, technetium,* rhenium, and polonium. Just as an example, we give in Table 1-2 the comparison of Mendeleev’s predictions of the properties of his hypothetical element "eka-silicon," with the actually observed properties of this element, which was found fifteen years later by a German chemist, Winkler, and given the name germanium.

Pretty good for a prediction at this stage in the development of chemistry!

By enumerating the elements from 1 (for hydrogen) and up as they come in the periodic system of elements, we obtain what is known as the atomic numbers of the elements.

*Technetium, an unstable element normally non-existent in nature, was produced only recently in atomic piles.

Table 1-2

Mendeleev’s prediction for eka-silicon (Es) (1871)

Winkler’s data for germanium (Ge) (Discovered in 1886)

Atomic weight will be about 72

Atomic weight is 72.6

Will be obtained from EsO2 or K2EsF6 by reduction with Na

Was obtained from K2GeF6 reduction with Na

Will be a dark gray metal with high melting point and density about 5.5

Is a gray metal with melting point 958° C and density 5.36

On heating, Es will form the oxide EsO2 with high melting point and density 4.7

Reacts with oxygen forming and density 4.7 GeO2 with melting point 1,100°C

The sulfide EsS2 will be insoluble in water but soluble in ammonium sulfide

GeS2 is insoluble in water but readily soluble in ammonium sulfide


Thus, the atomic number of carbon is 6, that of mercury is 80, and that of mendelevium, 101. The atomic numbers of the six noble gases that form important landmarks of chemical periodicity are: 2, 10, 18, 36, 54, and 86. It is convenient to represent the periodic system of elements by a three-dimensional spiral structure that is shown in Figure 1-2. The backbone of the structure is the column containing the noble gases running all the way from He down to Rn. The next column to the right contains the alkali metals, with hydrogen placed at the top because its chemical properties are similar to those of the alkali metals. To the left and around the corner from the noble gas column is the one containing the halogens. The first two periods, from He to F and from Ne to Cl, contain 8 elements each and fall neatly into this scheme, but the next period contains 18 elements and constitutes a problem. On the basis of chemical properties, there seems to be no doubt that the 3 elements that follow A (K, Ca, and Sc) must be placed under the 3 corresponding elements (Na, Mg, and Al) of the previous period and that those preceding Kr (As, Se, and Br) should be under those preceding A (P, S, and Cl), but we do not seem to have places for the remaining 11 elements Ti to Ge). To dispose of this difficulty, we place Ti and Ge, which both resemble Si, under that element and make an extra loop to accommodate the remaining 9 elements (V to Ga). The same situation arises in the next and in all of the following periods so that the extra loop perpetuates itself all the way to the end of the known sequence of elements. In the beginning of the fifth period, we encounter further trouble of the same kind and are forced to build another extra loop to accommodate 14 extra elements (Ce to Lu), known as the rare earths. The sixth and last period runs in the same way with most of the natural and artificial radioactive elements forming a loop under that formed by the rare earths.

Things become quite complicated, and Dmitri Ivanovich Mendeleev would probably be horrified by the looks of it, but that’s how it is. Nevertheless, in spite of the complexity of the diagram (which reflects the complexity of the internal structure of the atom), the periodic system of elements in Figure 1-2 gives a very good representation of the properties of the different elements.

Figure 1.2

Figure 1-2. The periodic system of the elements represented as a wound ribbon. The diagram on the next page shows the other side of the second loop. At present the ribbon is cut at atomic number 101 (mendelevium). An asterisk indicates that the element is unstable (radioactive), and an asterisk in parenthesis indicates the presence of a radioactive isotope in the normally stable element. The properties of the underlined elements were predicted by Mendeleev.

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