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
Figure 1.6 Ionic radii of the elements.
1.5.4 Chemical bonding
1.5.4.1 Covalent, ionic, and metal bonds
Except for the noble gases, atoms rarely exist independently; they are generally bound to other atoms in molecules, crystals, or ionic radicals. Atoms bind to one another through transfer or sharing of electrons, or through electrostatic forces arising from uneven distribution of charge in atoms and molecules. A bond that results from the transfer of electrons from one atom to another is known as an ionic bond, an example is the bond between Na and Cl in a halite crystal. In this case, the Na atom (the electropositive element) gives up an electron, becoming positively charged, to the Cl atom (the electronegative element), which becomes negatively charged. Electrostatic forces between the Na+ and the Cl– ions hold the ions in place in the crystal. When electrons are shared between atoms, such as in the H2O or CH4 molecules or the
Ideal covalent and ionic bonds represent the extremes of a spectrum: most bonds are neither wholly covalent nor wholly ionic. In these intermediate cases, the bonding electrons will spend most, but not all, of their time associated with one atom or another. Electronegativity is useful in describing the degree of ionicity of a bond: a bond is considered ionic when the difference in the electronegativity of the two atoms involved is greater than 2. In Figure 1.5, we see that metals tend to have low electronegativities while the nonmetals have high electronegativities. Thus, bonds between metals and nonmetals (e.g., NaCl) will be ionic while those between nonmetals (e.g., CO2) will be covalent, as will bonds between two like atoms (e.g., O2).
Another type of bond occurs in pure metal and metal alloy solids. In the metallic bond, valence electrons are not associated with any single atom or pair of atoms; rather, they are mobile and may be found at any site in the crystal lattice. Since metals rarely occur naturally at the surface of the Earth (they do occur in meteorites and the Earth's core), this type of bond is less important in geochemistry than other bonds.
Ionically bonded compounds tend to be hard, brittle, and highly soluble in water. Covalently bonded compounds tend to be good conductors of heat, but not of electricity. They are typically harder and less brittle than ionic solids but less soluble. In molecular solids, such as ice, atoms within the molecule are covalently bonded. The molecules themselves, which occupy the lattice points of the crystal, are bonded to each other through van der Waals and/or hydrogen bonds. Such solids are comparatively weak and soft and generally have low melting points.
Molecules in which electrons are unequally shared have an asymmetric distribution of charge and are termed polar. A good example is the hydrogen chloride molecule. The difference in electronegativity between hydrogen and chlorine is 0.9, so we can predict that the bonding electron will be shared but associated more with the Cl atom than with the H atom in HCl. Thus, the H atom will have a partial positive charge, and the Cl atom a partial negative charge. Such a molecule is said to be a dipole. The dipole moment, which is the product of one of the charges (the two charges are equal and opposite) times the distance between the charges, is a measure of the asymmetric distribution of charge. Dipole moment is usually expressed in Debye units (1 Debye = 3.3356 × 10–34 coulomb-meters).
1.5.4.2 Van der Waals interactions and hydrogen bonds
