Bonding Forces in Solids
Following section explains the bonding forces in solids in details.
- The interaction of electrons in neighboring atoms of a solid serves the very important function of holding the crystal together. For example, alkali halides such as NaCl are typified by ionic bonding.
- In the NaCl lattice, each Na atom is surrounded by six nearest neighbor CI atoms, and vice versa.
- Four of the nearest neighbors are evident in the two-dimensional representation shown in Fig.a.
The electronic structure of Na (Z = 11) is [Ne] 3s1, and CI (Z = 17) has the structure [Ne]3s23p5. In the lattice each Na atom gives up its outer 3s
electron to a CI atom, so that the crystal is made up of ions with the electronic structures of the inert atoms Ne and Ar (Ar has the electronic structure
- However, the ions have net electric charges after the electron exchange. The Na ion has a net positive charge, having lost an electron, and the Cl- ion has a net negative charge, having gained an electron.
- Each Na ion exerts an electrostatic attractive force upon its six Cl- neighbors, and vice versa.
- These coulombic forces pull the lattice together until a balance is reached with repulsive forces.
- A reasonably accurate calculation of the atomic spacing can be made by considering the ions as hard spheres being attracted together.
- In a metal atom the outer electronic shell is only partially filled, usually by no more than three electrons.
- We have already noted that the alkali metals (e.g., Na) have only one electron in the outer orbit. This electron is loosely bound and is given up easily in ion formation.
- This accounts for the great chemical activity in the alkali metals, as well as for their high electrical conductivity.
- In the metal the outer electron of each alkali atom is contributed to the crystal as a whole, so that the solid is made up of ions with closed shells immersed in a sea of free electrons.
- The forces holding the lattice together arise from an interaction between the positive ion cores and the surrounding free electrons.
- This is one type of metallic bonding.
- Obviously, there are complicated differences in the bonding forces for various metals, as evidenced by the wide range of melting temperatures (234 K for Hg, 3643 K for W).
- However, the metals have the sea of electrons in common, and these electrons are free to move about the crystal under the influence of an electric field.
A third type of bonding is exhibited by the diamond lattice semiconductors We recall that each atom in the Ge, Si, or C diamond lattice is surrounded
by four nearest neighbors, each with four electrons in the outer orbit.
- In these crystals each atom shares its valence electrons with its four neighbors (Fig.b).
- Bonding between nearest neighbor atoms is illustrated in the diamond lattice diagram.
- The bonding forces arise from a quantum mechanical interaction between the shared electrons.
- This is known as covalent bonding; each electron pair constitutes a covalent bond.
- In the sharingn process it is no longer relevant to ask which electron belongs to a particular atom—both belong to the bond.
- The two electrons are indistinguishable, except that they must have opposite spin to satisfy the Pauli exclusion principle.
- Covalent bonding is also found in certain molecules, such as H2.