Chapter 20: Metals and Their Compounds

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of the valence electrons by putting them lower in energy wells. The filled shells of the noble gases (helium, neon, argon, etc.) are particularly low energy arrangements of the electrons. These elements do not react with other elements because the reaction does not provide a means to an overall lowering of the electron energies. The effective oxidation state of the noble gases is identically zero. Many ions achieve this “noble gas electron configuration,” but of course no change has been made to the nucleus. In general, metallic elements near the left of the Periodic Table assume positive oxidation states in chemical reactions. Nonmetallic elements near the right side of the Periodic Table assume negative oxidation states. Near the middle of the Periodic Table things are somewhat ambiguous. Will an element like carbon (which has four valence electrons) lose its four valence electrons (like a metal), or gain four to fill its shell (like a nonmetal)? The answer depends on the specific compound that is being formed. An element in the middle, such as carbon, will do whichever best lowers its electron energy. Thus, for elements near the middle of the chart there is a tendency for atoms to assume different oxidation states in different chemical reactions, depending on the particular compound which is being formed. Although many elements can have more than one oxidation state (depending on the compound), the primary (likely) oxidation state is related to the atom’s position in the Periodic Table. All of the elements in column 1 (labeled IA in the Periodic Table) will have an oxidation state of +1 in ionic compounds. Elements in IIA have +2 oxidation states. What would you predict for column IIIA? Check your answer in Figure 20.4. Elements in columns VIA and VIIA of the Periodic Table tend to gain electrons to fill their valence shells and can be predicted with confidence to have primary oxidation states of –2 and –1, respectively. Elements in these columns sometimes have additional possible oxidation states in compounds. See sulfur and chlorine in Figure 20.4, for example. Elements in the B columns of the Periodic Table are also prone to having more than one possible oxidation state. However, we can use the additional information that these elements are all metals (and, therefore, tend to be electron donors) to predict that at least in some compounds these elements will have positive oxidation states that correspond directly to the number of valence electrons indicated by their column number. For example, chromium ( 24 Cr) in column VIB will in at least some cases have an oxidation state of +6. The rule is even less precise for the elements in the centermost columns of the Periodic Table (labeled VIII [neither A nor B]). Here we can only predict that, as metals, the elements will have positive oxidation states, but we must rely on experimental analysis of the individual compounds to determine the precise oxidation states of the constituent elements. Why do the regularities exist among the elements? They are based on ionization energies whose precise values can be traced back to the Wave Model and the Exclusion Principle. If only one electron in an atom has a low ionization energy, the oxidation state will be +1. If two electrons have low ionization energies, the oxidation state will be +2, and so on. If there are no electrons with low ionization energies, but there is a vacancy in an orbital deep in the well, the atom will try to fill the vacancy with an electron. This results in a –1 oxidation state. Compounds Between Metals and Nonmetals Imagine a metallic sodium ion, Na + , and a nonmetallic chloride ion, Cl – , that could form an ionic bond. But why should the pairing stop there? Could the Na + attract another negative ion and the Cl – attract another positive ion? Yes. Long chains of alternating Na + and Cl – ions form a sheet like a checkerboard (Fig. 20.5), and finally the checkerboards stack so that positive ions are always above negative ions and vice versa. The result is a crystal of salt based on the electrical attraction of the ions. Figure 20.5. Drawing of a small part of a single layer of a NaCl (sodium chloride) crystal. A salt is much different than a metal. Ionic bonding accounts for the characteristics of salts discussed below. 1. Electrically nonconducting. The electrons in both ions of a salt are all in stable, closed shells, and are difficult to remove from their respective ions. Because there are no mobile electrons, solid salts have no way to conduct electricity. However, if the salt is melted, whole ions may move about and conduct electricity, but they are so big and awkward they are not nearly as conductive as the electrons in a metal. Also, if the salt is dissolved in a liquid like water, the ions separate and are able to 186
conduct electricity, but in an inefficient manner. Materials that conduct electricity when dissolved in water are called electrolytes. 2. Transparent. The energy levels of most salts are spaced widely apart, so most salts will absorbs only photons with high energies. Such high-energy photons are in the invisible, ultraviolet part of the spectrum. Visible light is not absorbed by most salts—it passes through and the salts appear transparent. However, most of the salts of metals in columns IIIB through IIB absorb some visible light weakly; this gives them pale colors. There are also a few highly colored salts like potassium dichromate; most of them are composed of more than two elements (e.g., K 2 Cr 2 O 7 ). If a transparent salt crystal is ground to powder, the small crystals bend a light ray many times before it emerges from the powder. This means light will no longer pass straight through as it did in a single crystal, so light coming in from the side may bend and come out the front. The net result is that the powder appears white, even though it is basically transparent or slightly colored. A block of salt with many cracks or bubbles in it will also appear white. Many otherwise transparent minerals in nature appear opaque and white for this reason. 3. Brittle. Salt crystals break and shatter when hit with a hammer—they do not bend or flatten. Imagine sliding one layer of a NaCl crystal over the underlying layer (Fig. 20.6). When one sliding layer moves a short distance, all of the positive ions will be directly over other positive ions, and likewise for the negative ions. The repulsion is more than enough to force the two layers apart, which breaks (cleaves) the crystal. Brittleness, however, does not always dominate. Glass, most rocks, and objects made of fired clay can be considered super-cooled liquids, because under mild stress, only a few atoms at a time move and no breakage occurs. Over a long time many atoms will have moved, and the object will have a different shape. Thick rock layers bend, fold, and take on some surprisingly contorted shapes if they are not pushed too rapidly. If they cannot adapt quickly enough, they snap and cause earthquakes. Formulas and Names of Salts The charge in a crystal must be balanced (same total positive as negative charge). Even a relatively Figure 20.6. Force applied to layers of ions in a salt crystal will slide ions with the same charge over one another and split the crystal. small excess charge would cause a spark to jump from the crystal. This implies that there are nearly the same number of Na + as Cl – ions in a crystal of table salt. We can use subscripts to indicate how many ions there are in each crystal. For example, a small crystal containing 70 of each type of ion would be designated Na 70 Cl 70 . Because the proportion is always one-to-one, it is customary to write the simplest formula, NaCl. Other oneto-one compounds are LiF and CaO. Note that the metal is written first, the nonmetal last. Now consider Mg 2+ ions coming together with Cl – ions to form a crystal. Two Cl – ions are required to neutralize the +2 charge on the magnesium ion. The proportion is now one magnesium to two chlorides, and the formula is MgCl 2 . By knowing the oxidation numbers, we can determine the proper proportions and the proper subscripts in the formulas of compounds composed of a metal and a nonmetal. The pivotal concept is the cancellation of charges. You should be cautioned that this system works best for the combination of a metal with a nonmetal. It predicts just one compound—there may be others. If the oxidation states are known, the rule works for nonmetals also. For example, nitrogen in the unusual oxidation state of +4 can combine with oxygen in its –2 oxidation state. The resulting compound of two nonmetals has the formula NO 2 . However, it is technically incorrect to discuss such compounds in terms of oxida- 187
Page 1 and 2: 20. Metals and Their Compounds The
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Chapter 20: Metals and Their Compounds

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