The concepts of atomic radii, ionization energies, and electron affinities are properties of individual atoms. However, with the exception of the noble gases, none of the elements exist in nature as individual atoms. To get a broader understanding of the properties of elements, we must also examine periodic trends in properties that involve large collections of atoms.
The elements can be broadly groupted into the categories of metals, nonmetals, and metalloids. This classification is shown in Figure 7.9. Roughly three quarters of the elements are metals, situated in the left and middle portions of the table. The nonmetals are located at the top right corner, and the metalloids lie between the metals and nonmetals. Notice that hydrogen, which is located at the top left corner, is a nonmetal. It is for this reason that we set off hydrogen from the remaining group 1A elements. Some of the distinguishing properties of metals and nonmetals are summarized in Table 7.3.
Figure 7.9 The periodic table, showing metals, metalloids, and nonmetals.
The more an element exhibits the physical and chemical properties of metals, the greater its metallic character. Similarly, we can speak of the nonmetallic character of an element. As indicated in Figure 7.9, the metallic character generally increases as we proceed down a column of the periodic table and decreases as we proceed from left to right in a row. Let's now examine the close relationships that exist between electron configurations and the properties of metals, nonmetals, and metalloids.
Most metallic elements exhibit the shiny luster that we associate with metals (Figure 7.10). Metals conduct heat and electricity. They are malleable (can be pounded into thin sheets) and ductile (can be drawn into wire). All are solids at room temperature except mercury (melting point = -39°C), which is a liquid. Two melt at slightly above room temperature: cesium at 28.4°C and gallium at 29.8°C. At the other extreme, many metals melt at very high temperatures. For example, chromium melts at 1900°C.
Metals tend to have low ionization energies and are consequently oxidized (lose electrons) when they undergo chemical reaction. The relative ease of oxidation of common metals was discussed earlier, in Section 4.6. As we noted then, many metals are oxidized by a variety of common substances including O2 and acids.
Figure 7.11 shows the charges of some common ions. As we noted in Section 2.5, the charges of the alkali metals are always 1+ and those of the alkaline earth metals are always 2+ in their compounds. For each of these groups, the outer s electrons are easily lost, yielding a noble-gas electron configuration. The charges of the transition metal ions do not follow an obvious pattern. Many transition metal ions have 2+ charges, but 1+ and 3+ are also encountered. One of the characteristic features of the transition metals is their ability to form more than one positive ion. For example, iron may be 2+ in some compounds and 3+ in others. We will consider the electron configurations of these ions in Section 8.2.
Figure 7.11 Charges of some common ions found in ionic compounds. Notice that the steplike line that divides metals from nonmetals also separates cations from anions.
Compounds of metals with nonmetals tend to be ionic substances. For example, most metal oxides and halides are ionic solids. To illustrate, the reaction between nickel metal and oxygen produces nickel oxide, an ionic solid containing Ni2+ and O2– ions:
The oxides are particularly important because of the great abundance of oxygen in our environment.
Most metal oxides are basic oxides; those that dissolve in water react to form metal hydroxides, as in the following examples:
Metal oxides also demonstrate their basicity by reacting with acids to form salts and water, as illustrated in Figure 7.12.
(a) Write the chemical formula for aluminum oxide. (b) Would you expect this substance to be a solid, liquid, or gas at room temperature? (c) Write the balanced chemical equation for the reaction of aluminum oxide with nitric acid.
SOLUTION (a) In all its compounds, aluminum has a 3+ charge, Al3+; the oxide ion is O2–. Consequently, the formula of aluminum oxide is Al2O3.
(b) Because aluminum oxide is the oxide of a metal, we would expect it to be a solid. Indeed it is, and it has a very high melting point, 2072°C.
(c) Metal oxides generally react with acids to form salts and water. In this case the salt is aluminum nitrate, Al(NO3)3. The balanced equation is
Write the balanced chemical equation for the reaction between copper(II) oxide and sulfuric acid. Answer: CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l)
Nonmetals vary greatly in appearance (Figure 7.13). They are not lustrous and generally are poor conductors of heat and electricity. Their melting points are generally lower than those of metals (although diamond, a form of carbon, melts at 3570°C). Seven nonmetals exist under ordinary conditions as diatomic molecules. Included in this list are five gases (H2, N2, O2, F2, and Cl2), one liquid (Br2), and one volatile solid (I2). The remaining nonmetals are solids that can be hard like diamond or soft like sulfur.
Nonmetals, in reacting with metals, tend to gain electrons and become anions. For example, the reaction of aluminum with bromine produces aluminum bromide, an ionic compound containing the aluminum ion, Al3+, and the bromide ion, Br–:
The nonmetals commonly gain enough electrons to fill their outer p subshell completely, giving a noble-gas electron configuration (Figure 7.11).
Predict the formulas of compounds formed between (a) Ba and Te; (b) Ga and S.
SOLUTION (a) Notice the location of these elements in the periodic table. The charge on a barium ion is 2+; the charge on a tellurium ion (telluride) is 2-, like that of oxide and sulfide. Thus, the formula of the compound is BaTe.
(b) Gallium is in group 3A. Like the more familiar aluminum ion, gallium forms a 3+ ion. Sulfur forms a 2- ion (sulfide). Thus, the formula is Ga2S3.
Predict the formula of the compound formed by Rb and Se. Answer: Rb2Se
Compounds composed entirely of nonmetals are molecular substances. For example, the oxides, halides, and hydrides of the nonmetals are molecular substances that tend to be gases, liquids, or low-melting solids.
Most nonmetal oxides are acidic oxides; those that dissolve in water react to form acids, as in the following examples:
The reaction of carbon dioxide with water (Figure 7.14) accounts for the acidity of carbonated water and, to some extent, rainwater. Because sulfur is present in oil and coal, combustion of these common fuels produces sulfur dioxide and sulfur trioxide. These substances dissolve in water to produce acid rain, a major pollution problem in many parts of the world. The acidity of nonmetal oxides is also illustrated by the fact that they dissolve in basic solutions to form salts, as in the following examples:
Write the balanced chemical equations for the reactions of solid selenium dioxide with (a) water; (b) aqueous sodium hydroxide.
SOLUTION (a) Selenium dioxide is SeO2. Its reaction with water is like that of carbon dioxide (Equation 7.12):
(It doesn't matter that SeO2 is a solid and CO2 is a gas; the point is that both are soluble nonmetal oxides.)
(b) The reaction with sodium hydroxide is like the reaction summarized by Equation 7.14:
Write the balanced chemical equation for the reaction of solid tetraphosphorus hexoxide with water. Answer: P4O6(s) + 6H2O(l) 4H3PO3(aq)
Metalloids have properties intermediate between those of metals and nonmetals. They may have some characteristic metallic properties but lack others. For example, silicon looks like a metal (Figure 7.15), but it is brittle rather than malleable and is a much poorer conductor of heat and electricity than metals. Several of the metalloids, most notably silicon, are electrical semiconductors and are the principal elements used in the manufacture of integrated circuits and computer chips.