Our discussion has implied that the arrangement of elements in the periodic table followed knowledge of electron configuration. That implication is incorrect. In the late eighteenth and early nineteenth centuries, many elements were discovered and their properties described; those elements with similar properties were grouped in families and the families arranged in a periodic table. Figure 5.12 shows the location of these families in the periodic table.

FIGURE 5.12 The periodic table, showing the location of the representative, transition, and inner transition elements.

A. Families of Elements
One family of elements is the alkali metals: lithium, sodium, potassium, rubidium, cesium, and francium. These elements, found in column 1 of the periodic table, have a single valence electron. They are all soft, silvery gray solids with a clearly metallic luster. They are all very reactive, and their reactions with water are vigorous. These elements do not occur free in nature but are found combined with other elements, often with chlorine. The most common of these compounds is sodium chloride (table salt). Sodium chloride, NaCl, is the substance that makes seawater salty; it is also found in huge underground deposits (salt mines).

The alkaline earth metals, another family, are the elements in column 2 of the periodic table. Their atoms have two valence electrons. These metals are harder and stronger than the alkali metals. They, too, are found in combination with other elements. Although all alkaline earth metals have similar chemical reactivity, they have a wider range of reactivity than do the alkali metals. Beryllium and magnesium are unaffected by water, calcium reacts slowly with boiling water, and barium reacts violently with cold water. The elements of Group 7, with seven valence electrons, are known as the halogens. Chemically, these elements are very similar; physically, they are less alike. The lightest halogen, fluorine, is a pale yellow gas; iodine is a shiny, black solid. The heaviest halogen, astatine, is quite rare and is found in uranium ores. The total amount of astatine in the Earth's crust is probably less than 1 g. The longest-lived isotope of this element, astatine-210, has a half-life of only 8.3 hours. Several characteristic properties of halogens are shown in Table 5.6. Notice how these properties change as the atomic number increases.

The elements in column 8 are the noble gases. Recall from Chapter 3 that all these elements are monatomic gases and occur free and uncombined. They are all singularly unreactive, so much so that they were known earlier as the "inert gases." Only in 1960 were any of them shown to take part in chemical reactions. Even now, only krypton and xenon are known to form chemical compounds with other elements.

The elements in other columns of the table do not show striking similarities of properties. These columns are crossed by the stair-step line that separates metals from nonmetals, so they have nonmetals at the top and metals at the bottom of the column. For example, Group 4 is headed by the nonmetal carbon and has lead, a typical metal, as its heaviest member. Group 5 starts with the nonmetal nitrogen and ends with the metal bismuth.

B. Historical Development of the Periodic Table
By the middle of the nineteenth century, scientists knew the atomic weights of all the then-known elements and had observed that these weights progressed from 1 amu (hydrogen) to over 200 amu (lead). Efforts were made to find a systematic arrangement of all the elements that would place them in order of increasing atomic weight and in groups with similar properties. Such an arrangement, to be known as the periodic table, was conceived in 1869 both by Julius Lothar Meyer (1830-1395) in Germany and by Dmitri Mendeleev. (1834-1907) in Russia. Mendeleev is usually given most of the credit for constructing the periodic table because he not only described the ordering of the known elements but also used this arrangement to predict the existence of other elements not yet discovered.

Mendeleev's arrangement of the elements is shown in Figure 5.13. It resembles today's periodic table but has fewer columns and a different overall shape. Our current Group 8 is missing because the noble gases had not yet been discovered. The shorter columns of the d block are missing; instead, the transition elements are placed as subgroups in the long columns. Nevertheless, Mendeleev's table is the ancestor of our periodic table.

FIGURE 5.13 Mendeleev's periodic table (1871). The formulas of the oxide and the hydride of an element determined which column it would be in. Notice the spaces left empty but filled later as new elements were discovered. The atomic weights shown are those accepted in 1871; many have since been changed.

In making this table, Mendeleev followed most closely his aim to put similar elements in the same column. This aim had two important results. First, if the elements did not match, spaces were left and properties predicted for the elements yet to be discovered. Germanium, one of those for which a space was left, was discovered in 1886 by using Mendeleev's predicted properties as a guide. Second, if two neighboring elements were close in atomic weight and did not match their respective columns when placed in consecutive order, the order was reversed. Iodine and tellurium are two such elements. Although tellurium has a higher atomic weight, it is clearly not a halogen, but iodine is; thus, tellurium was placed in Group 6 and iodine in Group 7.

Mendeleev's arrangement of the elements held until about 1912 as an illustration of his statement of the periodic law: The properties of the elements are periodic functions of their atomic weights. In 1912 H. G. J. Moseley. (1887-1915), a student of Rutherford's, assigned to each element an atomic number that was equal to the charge on its nucleus and showed that this charge increased by one from one element to the next. Moseley's work suggested a restatement of the periodic law to read: the properties of the elements are periodic functions of their atomic numbers. This restatement resolved the tellurium-iodine problem as well as other similar problems. It recognized that atomic weights are based on isotopic distribution and that the real distinction between elements is the difference in the number of protons in their atoms. The periodic table we use today came into wide acceptance during the 1940s as the importance of electron configuration in predicting properties was recognized. Today we use the table as a summary of knowledge about the elements, whereas Mendeleev's table was held to be an interesting but not particularly useful artifact.