Types of Bonds in Crystals 晶体中键的类型

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Ionic Bonds

In some crystals , the atoms are present in a state where their electrons system is similar to that of a rare gas , so that their 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

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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 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 1s22s22p2 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 excited state 1s22s12p3 , 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 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 form a configuration of a regular tetrahedron.

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 accepter 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 neighbors. 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. Th …… 此处隐藏：4669字，全部文档内容请下载后查看。喜欢就下载吧 ……