Our discussion of atomic size, ionization energy, electron affinity, and metallic character gives some idea of the way the periodic table can be used to organize and remember facts. Not only do elements in a group possess general similarities, but there are also trends as we move through a group or from one group to another. In this section we will use the periodic table and our knowledge of electron configurations to examine the chemistry of the alkali metals (group 1A) and the alkaline earth metals (group 2A).
The alkali metals are soft metallic solids (Figure 7.16). All have characteristic metallic properties such as a silvery, metallic luster and high thermal and electrical conductivities. The name alkali comes from an Arabic word meaning "ashes." Many compounds of sodium and potassium, two alkali metals, were isolated from wood ashes by early chemists.
Sodium and potassium are among the most abundant elements in Earth's crust, in seawater, and in biological systems. Nearly everyone is aware of the presence of sodium ion in his or her body because of the use of salt in cooking and discusstions of the effects of dietary salt on blood pressure. Potassium is also prevalent in our bodies; a 140-pound person contains about 130 g of potassium, as K+ ion in intracellular fluids. Plants require potassium for growth and development (Figure 7.17).
Some of the physical and chemical properties of the alkali metals are given in Table 7.4. Notice that the elements have low densities and melting points and that these properties vary in a fairly regular way with increasing atomic number. We can also see some of the expected trends as we move down the group, such as increasing atomic radius and decreasing first ionization energy. For each row of the periodic table, the alkali metal has the lowest I1 value (Figure 7.6), which reflects the relative ease with which its outer s electron can be removed. As a result, the alkali metals are all very reactive, readily losing one electron to form ions with a 1+ charge:
(The symbol M represents any one of the alkali metals.) The alkali metals are the most active metals (Section 4.6) and thus exist in nature only as compounds. The metals can be obtained by passing an electric current through a molten salt, a process known as electrolysis. For example, sodium is prepared commercially by the electrolysis of molten NaCl. The electrical energy is used to remove electrons from Cl– ions and to force them onto Na+ ions:
We will discuss electrolysis in detail in Chapter 20. Not surprisingly, the chemistry of the alkali metals is dominated by the formation of 1+ cations (Equation 7.16). The metals combine directly with most nonmetals. For example, they react with hydrogen to form hydrides, with sulfur to form sulfides, and with chlorine to form chlorides:
In hydrides of the alkali metals (LiH, NaH, and so forth) hydrogen is present as H–, called the hydride ion. The hydride ion is distinct from the hydrogen ion, H+, formed when a hydrogen atom loses its electron.
The alkali metals react vigorously with water, producing hydrogen gas and solutions of alkali metal hydroxides:
These reactions are very exothermic. In many cases enough heat is generated to ignite the H2, producing a fire or explosion (Figure 7.18). This reaction is most violent in the case of the heavier members of the group, in keeping with their weaker hold on the single outer-shell electron.
The reactions between the alkali metals and oxygen are more complex. When oxygen reacts with metals, metal oxides, which contain the O2– ion, are usually formed. Indeed, lithium shows this reactivity:
When dissolved in water, Li2O and other soluble metal oxides react with water to form hydroxide ions, according to the net ionic equation
In contrast, the other alkali metals all react with oxygen to form metal peroxides, which contain the O22– ion. For example,
Potassium, rubidium, and cesium also form compounds that contain the O2– ion, called superoxides:
You should realize that the reactions shown in Equations 7.24 and 7.25 are somewhat surprising; in most cases, the reaction of oxygen with a metal forms the metal oxide.
As is evident from Equations 7.22 through 7.25, the alkali metals are extremely reactive toward water and oxygen. Because of this, the metals are usually stored under a hydrocarbon, such as kerosene or mineral oil.
Alkali metal salts and their aqueous solutions are colorless unless they contain a colored anion like the yellow CrO42–. Color is produced when an electron in an atom is excited from one energy level to another by visible radiation. Alkali metal ions, having lost their outermost electrons, have no electrons that can be excited by visible radiation.
When alkali metal compounds are placed in a flame, they emit characteristic colors (Figure 7.19). The alkali metal ions are reduced to gaseous metal atoms in the lower central region of the flame. The atoms are electronically excited by the high temperature of the flame; they then emit energy in the form of visible light as they return to the ground state. Sodium gives a yellow flame due to emission at 589 nm. This wavelength is produced when the excited state valence electron drops from the 3p subshell to the lower-energy 3s subshell. The characteristic yellow emission of sodium is the basis for sodium vapor lamps (Figure 7.20).
Write balanced equations that predict the reactions of cesium metal with (a) Cl2(g); (b) H2O(l); (c) H2(g).
SOLUTION We first recognize that cesium is an alkali metal. We therefore expect that its chemistry will be dominated by oxidation of the metal to Cs+ ions. Further, we recognize that Cs is far down the periodic table, which means it will be among the most active of all metals and will probably react with all three of the substances listed above.
By analogy to Equations 7.21, 7.22, and 7.19, we predict the occurrence of the following reactions:
In each case cesium forms a Cs+ ion in its compounds. The chloride (Cl–), hydroxide (OH–), and hydride (H–) ions are all 1- ions, so the final products have one-to-one stoichiometry with Cs+.
Write a balanced equation that predicts the products of the reaction between potassium metal and elemental sulfur. Answer: 2K(s) + S(s) K2S(s)
Like the alkali metals, the group 2A elements are all solids with typical metallic properties, some of which are listed in Table 7.5. Compared with the alkali metals, the alkaline earth metals are harder, are more dense, and melt at higher temperatures.
The first ionization energies of the alkaline earth elements are low, but not as low as those of the alkali metals. Consequently, the alkaline earths are less reactive than their alkali metal neighbors. As we have noted in Section 7.3, the ease with which the elements lose electrons decreases as we move across the periodic table from left to right and increases as we move down a group. Thus, beryllium and magnesium, the lightest members of the group, are the least reactive.
The trend of increasing reactivity within the group is shown by the behavior of the elements toward water. Beryllium does not react with water or steam, even when heated red-hot. Magnesium does not react with liquid water, but it does react with steam to form magnesium oxide and hydrogen:
Calcium and the elements below it react readily with water at room temperature (although more slowly than the alkali metals adjacent to them in the periodic table), as shown in Figure 7.22:
The two preceding reactions illustrate the dominant pattern in the reactivity of the alkaline earth elements—the tendency to lose their two outer s electrons and form 2+ ions. For example, magnesium reacts with chlorine at room temperature to form MgCl2, and it burns with dazzling brilliance in air to give MgO (Figure 3.6):
In the presence of O2, magnesium metal is protected from many chemicals by a thin surface coating of water-insoluble MgO. Thus, even though it is high in the activity series (Section 4.6), Mg can be incorporated into lightweight structural alloys used in, for example, automobile wheels. The heavier alkaline earth metals (Ca, Sr, and Ba) are even more reactive toward nonmetals than is magnesium, and they must be stored in such a way as to protect them from oxidation by O2 and H2O.
Because of their relatively high reactivity, the alkaline earth elements are invariably found in nature as compounds of the 2+ ions. Because the 2+ ions have a noble-gas electron configuration, they form colorless or white compounds unless they are combined with a colored anion. The heavier alkaline earth ions have characteristic flames; the calcium flame is brick red, strontium is crimson red, and barium is green.
Both magnesium and calcium are essential for living organisms (Figure 2.20). Calcium is particularly important for growth and maintenance of bones and teeth (Figure 7.23). In humans 99 percent of the calcium is found in the skeletal system.