Earth Materials. John O'Brien

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СКАЧАТЬ is much larger than the divalent sulfur cation (S⁺² = 0.37 Å) and the very small, highly charged hexavalent sulfur cation (S⁺⁶ = 0.12 Å). Keep in mind that the effective radius of a particular anion does vary somewhat. As we will see in the following sections, it depends on the environment in which bonding occurs, the number of nearest neighbors and the type of bond that forms.

2.3 CHEMICAL BONDS

      2.3.1 The basics

      Atoms in minerals, rocks, and other Earth materials are held together by forces or mechanisms called chemical bonds. The nature of these bonds strongly influences the properties and behavior of these materials. The nature of the bonds is, in turn, strongly influenced by the electron configuration of the elements that combine to produce the mineral, rock or other material.

      Five principle bond types and many hybrids occur in minerals. The three most common bond types are (1) ionic, (2) covalent, and (3) metallic. They can be modeled based on the behavior of valence electrons in the outer quantum levels of atoms. During bonding, valence electrons display varying tendencies to change position based on their periodic properties. In discussing chemical bonds, it is useful to divide elements into those that are metallic and those that are nonmetallic.

      Ionic bonds involve the linking together of metallic and nonmetallic elements, covalent bonds involve the linking of two nonmetallic elements, and metallic bonds involve the linking of two metallic elements. Hybrids between these bond types are common. Minerals with such hybrid or transitional bonds commonly possess combinations of features characteristic of each bond type. Other bond types include van der Waals and hydrogen bonds. Chemical bonding is a very complicated process; the models used below are simplifications designed to make this complex process easier to understand to a reasonable degree.

Another useful concept for understanding chemical bonds, developed originally by Linus Pauling (1929), is the concept of electronegativity (see Table 2.3). Electronegativity (En) is an empirical measure that expresses the tendency of an element to attract electrons when atoms bond. Highly electronegative elements (En > 3.0) have a strong tendency to attract electrons and become anions during bonding. Many column 16 (group VIA) and column 17 (group VIIA) elements are highly electronegative, requiring capture of two or one electrons, respectively, to achieve a stable electron configuration. Elements with low electronegativity (En < 1.5) are electropositive, metallic elements with a tendency to give up electrons to more electronegative elements during bonding to become positively charged cations. Highly electropositive elements include column 1 (group IA) and column 2 (group IIA) elements that tend to release one or two electrons, respectively, to achieve a stable electron configuration. Electronegativity is a very helpful concept in discussions of how atoms bond to produce larger molecules and minerals.

      2.3.2 Ionic (electrostatic) bonds

      When very metallic atoms bond with very nonmetallic atoms, an ionic bond, also called an electrostatic bond, is formed. Because the very metallic atoms (e.g., columns 1 and 2) are electropositive, they have a strong tendency to give up one or more electrons to achieve a stable configuration in their highest principal quantum level. In doing so, they become positively charged cations, whose charge is equal to the number of electrons each has lost. At the same time, very nonmetallic atoms (columns 16 and 17) are electronegative and have a strong tendency to gain one or more electrons in order to achieve a stable configuration in their highest principal quantum level. In doing so, they become negatively charged anions, with a charge equal to the number of electrons each has gained. When very metallic and very nonmetallic atoms bond, the metallic atoms give up or donate their valence electrons to the nonmetallic atoms that capture them. It is like a tug‐of‐war in which the electronegative side always wins the battle for electrons. In the electron exchange process, the atoms of both elements develop stable noble element electron configurations while becoming ions of opposite charge. Because particles of opposite charge attract, the cations and anions are held together by the electrostatic attraction between them that results from their opposite charges. Larger clusters of ions form as additional ions exchange electrons and are bonded and crystals begin to grow.

Figure 2.10 (a) Ionic bonding develops between highly electronegative anions and highly electropositive cations. When neutral sodium (Na⁰) atoms (red) release an electron to become cations (Na⁺¹) their ionic radius decreases. When neutral chlorine (Cl⁰) atoms (blue) capture an electron to become anions (Cl⁻¹) their ionic radius increases. (b) Ions of opposite charge attract to form crystals such as sodium chloride (NaCl).

The most frequently cited example of ionic bonding is the bonding between sodium (Na⁺¹) and chloride (Cl⁻¹) ions in the mineral halite (NaCl) (Figure 2.10). As a column 1 (group IA) element, sodium is very metallic and electropositive, with a rather low electronegativity (0.93). Sodium has a strong tendency to give up one electron to achieve a stable electron configuration. On the other hand, chlorine, as a column 17 (group VIIA) element, is very nonmetallic, has a strong affinity for electrons and has a high electronegativity (3.5). It has a strong tendency to gain one electron to achieve a stable electron configuration. When sodium and chlorine atoms bond, the sodium atoms release one electron to become smaller sodium cations (Na⁺¹) with the “stable octet” electron configuration (neon), while the chlorine atoms capture one electron to become larger chloride anions (Cl⁻¹) with “stable octet” electron configurations. As we shall see, the “exchange” is incomplete. The two atoms are then joined together by the electrostatic attraction between particles of opposite charge to form the compound NaCl. In macroscopic mineral specimens of halite, many millions of sodium and chloride ions are bonded together, each by the electrostatic or ionic bond described above. Note that the numbers of chloride anions and sodium cations must be equal if the electric charges are to be balanced so that the mineral is electrically neutral. Other group IA (1) and group VIIA (17) elements bond ionically to produce minerals such as sylvite (KCl).

      Ionic bonds also form when group IIA and group VIA elements combine. In the mineral periclase (MgO), magnesium (Mg⁺²) and oxygen (O⁻²) ions are bonded together to form MgO. In this case, electropositive, metallic magnesium atoms from group IIA tend to donate two valence electrons to become stable, smaller divalent magnesium cations (Mg⁺²) while highly electronegative, nonmetallic oxygen atoms from group VIA capture two valence electrons to become stable, larger divalent oxygen anions (O⁻²). The two oppositely charged ions are then held together by virtue of their opposite charges by an electrostatic or ionic bond. Once again, the number of magnesium cations (Mg⁺²) and oxygen anions (O⁻²) in periclase (MgO) must be the same if electrical neutrality is to be conserved. A slightly more complicated example of ionic bonding involves the formation of the mineral fluorite (CaF₂). In this case, electropositive, metallic calcium atoms from class IIA release two electrons to become stable divalent cations (Ca⁺²). At the same time, two nonmetallic, strongly electronegative fluorine atoms from class VIIA each accept one of these electrons to become stable univalent anions (F⁻¹). Pairs of F⁻¹ anions bond to each Ca⁺² cation to form ionic bonds in electrically neutral fluorite (CaF₂).

In ideal ionic bonds, ions can be modeled as spheres of specific ionic radii in contact with one another (Figure 2.10), as though they were ping‐pong balls or marbles in contact with each other. This approximates real СКАЧАТЬ