CoreChem:4.2 The Periodic Classification of the Elements
From ChemEd Collaborative
4.2 THE PERIODIC CLASSIFICATION OF THE ELEMENTS
The similarities among macroscopic properties within each of the chemical families just described lead one to expect microscopic similarities as well. Atoms of sodium ought to be similar in some way to atoms of lithium, potassium, and the other alkali metals. This could account for the related chemical reactivities and analogous compounds of these elements.
According to Dalton’s atomic theory, different kinds of atoms may be distinguished by their relative masses (atomic weights). Therefore it seems reasonable to expect some correlation between this microscopic property and macroscopic chemical behavior. You can see that such a relationship exist by listing symbols for the first dozen elements in order of increasing relative mass. Obtaining atomic weights from Table 2.2, we have
Elements which belong to families we have already discussed are indicated by shading around their symbols. The second, third, and forth elements on the list (He, Li, and Be) are a noble gas, an alkali metal, and an alkaline-earth metal, respectively. Exactly the same sequence is repeated eight elements later (Ne, Na, and Mg), but this time a halogen (F) precedes the noble gas. If a list were made of all elements, we would find the sequence halogen, noble gas, alkali metal, and alkaline-earth metal several more times.
The Periodic Table
In 1871 the Russian chemist Dmitri Ivanovich Mendeleev (1834 to 1907) proposed the periodic law. This law states that when the elements are listed in order of increasing atomic weights, their properties vary periodically. That is, similar elements do not have similar atomic weights. Rather, as we go down a list of elements in order of atomic weights, corresponding properties are observed at regular intervals. To emphasize this periodic repetition of similar properties, Mendeleev arranged the symbols and atomic weights of the elements in the table shown in Fig. 4.1. Each vertical column of this periodic table contains a group or family of related elements. The alkali metals are in group I (Gruppe I), alkaline earths in group II, chalcogens in group VI, and halogens in group VII. Mendeleev was not quite sure where to put the coinage metals, and so they appear twice. Each time, however, copper, silver, and gold are arranged in a vertical column. Although the noble gases were discovered nearly a quarter century after Mendeleev’s first periodic table was published, we have included them in Fig. 4.1 to indicate that they, too, fit the periodic arrangement. In constructing his table, Mendeleev found that sometimes there were not enough elements to fill all the available spaces in each horizontal row or period. When this was true, he assumed that eventually someone would discover the element or elements needed to complete a period.
Figure 4.1 Mendeleev’s periodic table, redrawn from “Annalen der Chemie,” supplemental volume 8, 1872. The German words Gruppe and Reihen indicate, respectively, the groups and rows (or periods) in the table. Mendeleev also used the European convention of a comma instead of a period for the decimal and J instead I for iodine. The noble gases had not yet been discovered when Mendeleev devised the periodic table, but they have been included here (in color) for completeness.
TABELLE II
TABLE 4.1 Comparison of Mendeleev’s Predictions with the Observed Properties of the Element Scandium.
| Properties Predicted for Ekaboron (Eb)* by Mendeleev 1872 | Properties Found for Scandium after its Discovery in 1879 | |
| Atomic weight | 44 | 44† |
| Formula of oxide | Eb2O3 | Sc2O3 |
| Density of oxide | 3.5 | 3.86 |
| Acidity of oxide | Greater than MgO | Greater than MgO |
| Formula of chloride | EbCl3 | ScCl3 |
| Boiling point of chloride | Higher than for | Higher than for |
| Color of compounds | Colorless | Colorless |
* Mendeleev used the name ''eka''boron because the blank space into which the element should fit was ''below'' boron in his periodic table.
† The modern value of the atomic weight of scandium is 44.96.
Therefore he left blank spaces for undiscovered elements and predicted their properties by averaging the characteristics of other elements in the same group.
As an example of this process, look at the fourth numbered row (Reihen) in Fig. 4.1. Scandium (Sc) was unknown in 1872; so titanium (Ti) followed calcium (Ca) in order of atomic weights. This would have placed titanium below boron (B) in group III, but Mendeleev knew that the most common oxide of titanium, TiO2, had a formula similar to an oxide of carbon CO2, rather than of boron, B2O3. Therefore he placed titanium below carbon in group IV. He proposed that an undiscovered element, ekaboron, would eventually be found to fit below boron. (The prefix eka means “below.”) Properties predicted for ekaboron are shown in Table 4.1. They agreed remarkably with those measured experimentally for scandium when it was discovered 7 years later. This agreement was convincing evidence that a periodic table is a good way to summarize a great many macroscopic, experimental facts.
The modern periodic table inside the front cover of this book differs in some ways from Mendeleev’s original version. It contains more than 40 additional elements, and its rows are longer instead of being squeezed under one another in staggered columns. (Mendeleev’s fourth and fifth rows are both contained in the fourth period of the modern table, for example.) The extremely important idea of vertical groups of related elements is still retained, as are Mendeleev’s group numbers. The latter appear as roman numerals at the top of each column in the modern table.
Valence
Perhaps the most important function of the periodic table is that it helps us to predict the chemical formulas of commonly occurring compounds. At the top of each group, Mendeleev provided a general formula for oxides of the elements in the group. (See Fig. 4.1). The heading R2O above group I, for example, means that we can expect to find compounds such as H2O, Li2O, Na2O etc. Similarly, the general formula RH3 above group V suggests that the compounds NH3, PH3,VH3, and AsH3(among others) should exist. To provide a basis for checking this prediction, formulas are shown in Table 4.2 for compounds in which H, O, or Cl is combined with each of the first two dozen elements (in order of atomic weights). Even among groups of elements whose descriptive chemistry we have not discussed, you can easily confirm that most of the predicted formulas correspond to compounds which actually exist. Conversely, more than 40 percent of the formulas for known O compounds agree with Mendeleev’s general formulas. (These are shaded in color in Table 4.2.) The periodic repetition of similar formulas is even more pronounced in the case of Cl compounds. This is evident when a list is made of subscripts for Cl in combination with each of the first 24 elements. Consulting Table 4.2, we find HCl (subscript 1), no compound with He (subscript 0), LiCl (subscript 1),and so on.
With only the two exceptions indicated in italics, at least one formula for a compound of each element fits a sequence of subscripts which fluctuate regularly from 0 up to 4 and back to0 again. (The unusual behavior of K and Ar will be discussed a bit later.) The number of Cl atoms which combines with one atom of each other element varies quite regularly as the atomic weight of the other element increases.
The experimentally determined formulas in Table 4.2 and the general formulas in Mendeleev’s periodic table both imply that each element has a characteristic chemical combining capacity. This capacity is called valence, and it varies periodically with increasing atomic weight. The noble gases all have valences of 0 because they almost never combine with any other element. H and Cl both have the same valence.
TABLE 4.2 Molecular Formulas for Hydrogen, Oxygen, and Chlorine Compounds of the First Twenty-Four Elements in Order of Atomic Weight.*
* For each element compounds are listed in order of decreasing stability. In some cases additional compounds are known, but these are relatively unstable.
† A great many stable compounds of carbon and hydrogen are known, but space limitations prevent listing all of them.
They combine with each other in a 1:1 ratio to form HCl, each combines with Li in the same 1:1 ratio (LiH and LiCl), each combines with Be in the same ratio (BeH2, BeCl2), and so on. Because H and Cl have the same valence, we can predict that a large number of H compounds will have formulas identical to those of Cl compounds, except, of course, that the symbol H would replace the symbol Cl. The correctness of this prediction can be verified by studying the formulas surrounded by gray shading in Table 4.2
The combining capacity, or valence, of O is apparently twice that of H or Cl. Two H atoms combine with one O atom in H2O So do two Cl atoms or two Li atoms (Cl2O and Li2O). The number of atoms combining with a single O atom is usually twice as great as the number which combined with a single H or Cl atom. (Again, consulting the gray shaded formulas in Table 4.2 will confirm this statement.)
After careful study of the formulas in the table, it is also possible to conclude that none of the elements (except the unreactive noble gases) have smaller valences than H or Cl. Hence we assign a valence of 1 to H and to Cl. The valence of O is twice as great, and so we assign a value of 2.
EXAMPLE 4.1 Use the data in Table 4.2 to predict what formula would be expected for a compound containing (a) sodium and fluorine; (b) calcium and fluorine.
Solution
a) From the table we can obtain the following formulas for the most common sodium compounds:
- NaHNa2ONaCl
All of these would imply that sodium has a valence of 1. For fluorine compounds we have
- HFOF2ClF
which imply that fluorine also has a valence of 1. Therefore the formula is probably
- NaF
b) We already know that the valence of fluorine is 1. For calcium the formulas
- CaH2CaOCaCl2
argue in favor of a valence of 2. Therefore the formula is most likely
- CaF2
In some cases one element can combine in more than one way with another. For example, you have already encountered the compounds HgBr2 and Hg2Br2. There are many other examples of such variable valence in Table 4.2. Nevertheless in its most common compounds, each element usually exhibits one characteristic valence, no matter what its partner is. Therefore it is possible to use that valence to predict formulas. Variable valence of an element may be looked upon as an exception to the rule of a specific combining capacity for each element.
The experimental observation that a given element usually has a specific valence can be explained if we assume that each of its atoms has a fixed number of valence sites. One of these sites would be required to connect with one site on another atom. In other words, a noble-gas atom such as Ar or Ne would not have any combining sites, H and Cl atoms would have one valence site each, an O atom would have two, and so on. Variable valence must involve atoms in which some valence sites are more readily used than others. In the case of the F compounds of Cl (ClF, ClF3, ClF5), for example, the formulas imply that at least five valence sites are available on Cl. Only one of these is used in ClF and in most of the chlorine compounds of Table 4.2. The others are apparently less readily available.
Mendeleev’s inclusion of general formulas above the columns of his periodic table indicates that the table may be used to predict valences of the elements and formulas for their compounds. Two general rules may be followed:
1 In periodic groups I to IV, the group number is the most common valence.
2 In periodic groups V to VII, the most common valence is equal to 8 minus the group number, or to the group number itself.
For groups V to VII, the group number gives the valence only when the element in question is combined with oxygen, fluorine, or perhaps one of the other halogens. Otherwise 8 minus the group number is the rule.
EXAMPLE 4.2 Use the modern periodic table inside the front cover of this book to predict the formulas of compounds formed from (a) aluminum and chlorine; (b) phosphorus and chlorine. Use Table 4.2 to verify your prediction.
Solution
a) Aluminum is in group III and so rule 1 predicts a valence of 3. Chlorine is in group VII and is not combined with oxygen or fluorine, and so its valence is 8 – 7 = 1 by rule 2. Each aluminum has three valence sites, while each chlorine has only one, and so it requires three chlorine atoms to satisfy one aluminum, and the formula is AlCl3.
b) Again chlorine has a valence of 1. Phosphorus is in group V and might have a valence of 5 or of 8 – 5 = 3. Therefore we predict formulas PCl5 or PCl3. Note: All three predicted formulas appear in Table 4.2.
Exceptions to the Periodic Law
In the process of constructing the first periodic table, Mendeleev encountered several situations where the properties of elements were incompatible with the positions they would be forced to occupy in order of increasing atomic weight. In such a case, Mendeleev chose to emphasize the properties, because in the 1870s it was difficult to determine atomic weights accurately. He assumed that some atomic weights were in error and that ordering of elements ought to be changed to agree with chemical behavior. We pointed out a problem of this type in the preceding section. Mendeleev did not have to contend with it because the noble gases had not been discovered in 1872, but it illustrates the difficulty nicely. There was a break in the regular sequence of valences of the first 24 elements when we came to K and Ar. The alkali metal has a smaller atomic weight than the noble gas and appears before the noble gas in Table 4.2. All other alkali metals immediately follow noble gases (they have slightly larger atomic weights). Unless we make an exception to the order of increasing atomic weight for Ar and K, the periodic table would contain a strange anomaly. One of the elements in the vertical column of noble gases would be the extremely reactive K. Likewise, the group of alkali metals would contain Ar, which is not a metal and is very unreactive. Mendeleev’s assumption that more accurate atomic weight determinations would eliminate situations such as we have just described has turned out to be incorrect. The atomic weights in Table 4.2 are modern, highly accurate values, but they still predict the wrong order for Ar and K. The same problem occurs in the case of Co and Ni and of Te and I. Apparently atomic weight, although related to chemical behavior, is not as fundamental as Mendeleev and other early developers of the periodic table thought.
Implications of Periodicity for Atomic Theory
The concept of valence implies that atoms of each element have a characteristic number of sites by which they can be connected to atoms of other elements. The number of valence sites repeats periodically as atomic weight increases, and occasionally even this regular repetition is imperfect. Atoms of similar atomic weight often have quite different properties, while some which differ widely in relative mass behave almost the same. Dalton’s atomic theory considers atoms to be indestructible spheres whose most important property is mass. This is clearly inadequate to account for the macroscopic observations described in this and the preceding section. In order to continue using the atomic theory, we must attribute some underlying structure to atoms. If both valence and atomic weight are determined by that structure, we should be able to account for the close but imperfect relationship between these two properties. The next section will describe some of the experiments which led to current theories about just what this atomic structure is like.
