Unit 1 Types of Bonds in Crystals
Ionic Bonds
In some crystals, the atoms are present in a state where their electron system is similar to that of a rare gas, so that outer shell has either lost excess electrons or has been filled with a total of eight electrons, i.e. , completed. Since the electric neutrality must be conserved, the crystal always simultaneously contains atoms that donate electrons and atoms that accept electrons. The former form positively charged cations, while the latter form negatively charged anions, and the electric charges of ions are integral multiples of the electron charges.
The cohesive forces in these crystals are electrostatic forces acting between the ions. This type of bond is called an ionic bond and the crystals are called ionic crystals. As the electric field of ions is spherically symmetrical, the ionic bonding is isotropic, i.e., the bonds do not have directional character and every ion attempts to be surrounded by the maximum possible number of ions of the opposite charge, so that the bonds are not saturated. Halide salts of alkali metals represent typical ionic crystals; this is so because alkali metals have only one electron in the outer shell, while halides lack exactly one electron for completion of their outer shell to eight electrons.
This idea is identical with the concept of valency except that the crystal is not considered as a compound of molecules, but rather as a unified structure for which the chemical formula has the meaning of the ratio of the elements and the geometric arrangement is an indispensable part of the description of the substance. The chemical formula, e.g. , NaCl, does not denote a molecular structural unit here, because every ion in the crystal interacts with several closest neighbours with the opposite sign, so that, for example, each Na+ ion in NaCl is surrounded by six equivalent nearest Cl- ions, and vice versa.
Covalent Bonds
An exact quantum mechanical calculation for the hydrogen molecule model, carried out by Heitler and London (1927), revealed that there exist two possible lowest energy states of the hydrogen molecule, composed of the original single-atomic states, and that the lower energy corresponds to a singlet state in which orientations of the spins of electrons are antiparallel. The energy difference between the two states and the consequent forces, called exchange forces, depend on the overlap of the wave functions of the electrons, which become common for both atoms. Such a bonding is called homeopolar or atomic.
The pairing of electrons in states in which the electrons, according to Pauli principle, differ only in spin orientation is also characteristic for covalent bonds between atoms with more complex electron structures. The bond is again created by the overlap of the single electron wave functions of atomic orbitals, which combine into the wave function of the common state. The main characteristics of covalent bonds are their saturation and mutual orientation of the bonds when there are several on the given atom; this is always the case except for atomic pairs. Saturation is a consequence of the Pauli principle; every bond contains exactly two electrons. Formally, the bond is similar to an ionic bond with the difference that the electrons are not transferred between atoms but become common property. For electrons in the p- and d- states, the degree of overlap, and thus also the covalent bond, depends not only on the interatomic distances but also on the mutual orientation of the directions of lines connecting the various atoms.
Bond hybridization
A future important property of covalent bonds is hybridization of the atomic orbitals, leading to variation of the valency of a given element in various compounds. The best example of this phenomenon is carbon, the ground state of which has the electron configuration (1s2, 2s2, 2p2) with only two p-electrons unpaired, which results in valency of two as, for example, in CO. However, it is known that carbon is usually present in compounds as a tetravalent element. Pauling (1931) explained this phenomenon through orbital hybridization, where carbon is present in the compound in the excited state (1s2, 2s1, 2p3), so that it has one unpaired s-electron and three unpaired p-electrons, i.e., a total of four electrons that can enter into covalent bonds. From a quantum mechanical point of view, these electrons cannot be considered separately but must be considered as equivalent particles in a single common state. From this also follows that the four bonds, available on carbon in this state, are completely equivalent. As a consequence, the atoms bonded to tetravalent carbon from a configuration of a regular tetrahedron. According to Kimball, the spatial orientations of the covalent bonds have the following shapes for various combinations of atomic orbitals:(table 1, P2).
Metallic bonds
In metallic crystals, many of the valence electrons are present in states where they are not localized close to the atoms but can move freely over the crystal, which is reflected in a high conductivity. The structure of metals can thus be conceived of as being composed of positively charged ions immersed in an electron gas. The term "gas" is not used accidentally here, as the free electrons actually behave statistically like a gas. The bonding between the positive ions occurs through these free electrons. Thus, an ionic bond can be described as containing an electron localized close to an acceptor atom, while the electron in a covalent bond is located between two atoms, and in a metallic bond the free electrons are completely delocalized. Metallic bonding, similarly to ionic, is characterized by isotropy and the consequent large coordination numbers. The ions in a metal tend to be surrounded by the maximum possible number of neighbours. In contrast to ionic bonds, metallic bonds do not require a balance of the electric charge between the elements; the electrostatic equilibrium is between metal ions and electron gas. This is why different metal elements can mix in crystals in practically arbitrary ratios to form alloys.
Van der Waals bonds
Van der Waals bonds are familiar from the kinetic theory of gases as the forces responsible for the deviation of the behaviour of a gas from that of an ideal gas. These are attractive forces between molecules with saturated bonds. Of all the forces in crystals, they are the weakest and mostly responsible for the formation of layered structures with weak cohesion between the layers. Van der Waals forces are present in molecular crystals, in which the molecules constitute the basic structural units. The forces which make the molecules stable are either of ionic or of covalent character while the molecules themselves are bound into crystal by van der Waals forces.
Selected from "Structure and Properties of Ceramics", A. Koller, Elsevier Publ., 1994
Words and Expressions
cation n.阳离子
anion n.阴离子
cohesive a.内聚（力）的
spherically ad.球（形）的
isotropic a.各向同性的
halide n.卤化物
indispensable a.必不可少的
vice versa 反之亦然
singlet n.零自旋能级/自旋单态
antiparallel a.反平行的
overlap n.vt.重叠
wave function 波函数
homopolar a.同极化的
mutual orientation 相互取向
hybridization n.杂化
configuration n.构型/电子排布
tetravalent n.四价的
orbital hybridization 轨道杂化
tetrahedron n.四面体
spatial orientation 空间取向
trigonal a.三方的
pyramidal a.四方锥的
tetragonal a.四方的
bipyramidal a.双四方锥的
octahedral a.八面体的
prismatic a.斜方晶系的
Van der Waals forces 范德瓦尔斯力/范德华力
coordination number 配位数
