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It is only the most loosely bound electrons, those in the outermost shells, that participate in chemical bonding, so elements sharing a similar outermost electronic configuration tend to behave similarly. Elements within the same column of the periodic table, or group, share outer electronic configurations and hence behave in a similar manner. Thus, the elements of Group 1, the alkalis, all have one electron in the outermost s orbital, and behave in a similar manner. The Group 18 elements, the noble, or rare, gases, all have a filled p subshell, and behave similarly.

Let's now consider several concepts that are useful in describing the behavior of atoms and elements: ionization potential, electron affinity, and electronegativity. The first ionization potential of an atom is the energy required to remove (i.e., move an infinite distance away) the least tightly bound electron. This is energy gained by the electron in reactions such as:

(1.1)

The first ionization potential of the elements is illustrated in Figure 1.4. The Second Ionization Potential is the energy required to remove a second electron, etc. The electron affinity is the energy given up in reactions such as:

(1.2)

Electronegativity is another parameter that is often used to characterize the behavior of the elements. It is a relative, unitless quantity determined from the differences in bond energy between an A–B molecule and the mean energies of A–A and B–B molecules. Electronegativity quantifies the tendency of an element to attract a shared electron when bonded to another element. For example, F has a higher electronegativity than H (the values are 3.8 and 2.5, respectively), thus the bonding electron in hydrogen fluoride, HF, is more likely to be found in the vicinity of F than of H. It is also useful in characterizing the nature of chemical bonds between elements, as we shall see in a subsequent section. Electronegativities of the elements are shown in Figure 1.5.

In general, first ionization potential, electron affinity, and electronegativities increase from left to right across the periodic table, and to a lesser degree from bottom to top. This reflects the shielding of outer electrons, particularly those in s orbitals, by inner electrons, particularly those in p orbitals, from the charge of the nucleus. Thus the outer 3s electron of neutral sodium is effectively shielded from the nucleus and is quite easily removed. On the other hand, the 2p orbitals of oxygen are not very effectively shielded, and it readily accepts two additional electrons. With the addition of these two electrons, the 2p orbital is filled and the 3s orbital effectively shielded, so there is no tendency to add a third electron. With the outer p (and s) orbitals filled, a particularly stable configuration is reached. Thus, Ne and other noble gases have little tendency to either add or give up an electron.

Figure 1.4 First ionization potential of the elements.

Figure 1.5 Electronegativities of the elements. Nonmetals are characterized by high electronegativity, metals by low electronegativity. Metalloids have intermediate values.

Metallic elements have electronegativities generally ≤ 1.9 and are said to be “electropositive”. They tend to form positively charged ions, called cations, by giving up electrons. Elements with electronegativities ≥2.5 are nonmetals and tend to form negatively charged ions, called anions, by acquiring additional elections. Those with electronegativities in the range of >1.8 and <2.2 are called metalloids or semi-metals and form either type of ion.

The number of electrons that an element will either give up or accept is known as its valence. For elements in the wings of the periodic table (i.e., all except the transition metals), valence is easily determined simply by counting how far the element is horizontally displaced from Group 18 in the periodic table. For Group 18, this is 0, so these elements, the noble gases, have 0 valence. For Group 1 it is 1, so these elements have valence of +1; for Group 17 it is –1, so these elements have valence of –1, etc. Valence of the transition metals is not so simply determined, and these elements can have more than one valence state. Most, however, have valence of 2 or 3, though some, such as U, can have valences as high as 6.

A final characteristic that is important in controlling chemical properties is ionic radius. This is deduced from bond length when the atom is bonded to one or more other atoms. Positively charged atoms, or cations, have smaller ionic radii than do negatively charged atoms, or anions. Also, ionic radius decreases as charge increases for cations. This decrease is due both to loss of outer electrons and to shrinking of the orbits of the remaining electrons. The latter occurs because the charge of the nucleus is shared by fewer electrons and hence has a greater attractive force on each. In addition, ionic radius increases downward in each group in the periodic table, both because of addition of electrons to outer shells and because these outer electrons are increasingly shielded from the nuclear charge by the inner ones. Ionic radius is important in determining geochemical properties such as substitution in solids, solubility, and diffusion rates. Large ions are surrounded, or coordinated, by a greater number of oppositely charged ions than do smaller ones. The ionic radii of the elements are illustrated in Figure 1.6.

We can now summarize a few of the more important chemical properties of the various groups in the periodic table. Group 18 does not participate in chemical bonding in nature, hence the term noble gases. Group 1 elements, the alkalis, readily lose an electron and hence are highly reactive. They tend to form ionic bonds rather than covalent ones and hence weaker bonds to other elements and to be quite soluble in aqueous solutions and can be easily leached from minerals. Because they have only a +1 charge, their ionic radii tend to be larger than those of other cations. Group 2 elements, the alkaline earths, have these same characteristics, but somewhat moderated. Group 17 elements, the halogens, are highly electronegative and readily accept an electron, are highly reactive, form ionic bonds, and are quite soluble. Their ionic radii tend to be larger than more highly charged anions. Elements of Groups 13–16 tend to form bonds that are predominantly covalent. As a result, they tend to be less reactive and less soluble (except where they form soluble radicals, such as ) than Group 1, 2, and 17 elements. Finally, the transition metals are a varied lot. Many form strong bonds (generally with O in nature) and can have quite variable solubility (which, as we shall see in Chapter 6, depends on their valence state). Some, the noble metals (Ru, Rh, Pd, Os, Ir, Pt) in particular, tend to be very unreactive. The rare earths are of interest because all have two electrons in the 6s outer orbital and vary only by the number of electrons in the 4f shell. Because they almost always have the same valence (+3), bonding behavior is quite similar. As Goldschmidt found, they vary systematically in ionic radius, which results in a systematic variance in geochemical behavior, as we will see in Chapter 7.

Figure 1.6 Ionic radii of the elements.