Interatomic Bonding in the Solid-State

The term solid-state came from the fact that the transistor technology used solid-state materials instead of the gaseous state employed in vacuum tubes. The prominent characteristics of solids are exact size and shape and much greater corporeality than gases. The symmetry of the arrangement of atoms in solids has stimulated and allowed rapid progress in the field of solid-state electronics.

 

A quartz crystal, composed of silicon and oxygen (SiO2). 

 

Atomic Bonding

Most materials used in electronics nowadays are solid. The Electronics Engineer needs to know the attractions that hold the atoms together since the materials’ engineering properties depend on the interatomic forces present.

The electronic structure of the atoms originates from the interatomic attractions. The eight-electron configuration in the valence subshells s and p is stable, and the atoms with less than eight electrons tend to react, looking for more stable compounds. An exception is the K-shell, which is stable with two electrons.

The atoms may achieve bonding by filling their outer s and p levels receiving extra electrons, giving up electrons, or sharing electrons.

The first two of these procedures produce ions with a net negative or positive charge to provide coulombic attraction to other ions of opposite charge – the ionic bond. The third procedure requires a close association between atoms to share electrons – the covalent bond.

The bonds between adjacent atoms resulting from the transfer or sharing of outer electrons are relatively strong.

In addition to ionic and covalent bonds, there is a third type of interatomic force capable of holding atoms together: the metallic bond.

These types are the primary interatomic bonds.

The Van der Waals bonds are secondary bonds originating from a distinct mechanism and are relatively weaker. If it weren’t that they are sometimes the only forces present, they could go unnoticed. The origin of van der Waals forces between atoms is quantum mechanical and beyond this article’s scope.

 

The Ionic Bond

The ionic bond is the easiest one to describe. When a material has several types of atoms, one atom may donate its valence electrons to another, filling the second atom’s valence shell.

Both atoms now have filled – or emptied – valence shells and have acquired an electrical charge, behaving as ions. The mutual attraction of the negative and positive charges creates the ionic bond.

Figure 1 shows sodium chloride production (NaCl or table salt) by the attraction between sodium and chloride ions.

Figure 1. Ionic bond in NaCl.

 

The Covalent Bond

Sometimes atoms can complete the eight electrons in the valence shell by sharing them with adjacent atoms.

Covalent bonds provide great attraction forces between atoms. An example is a diamond, the hardest material found in nature, which is pure carbon. Each carbon atom has four electrons in its valence shell, shared with four adjacent atoms to form a three-dimensional lattice entirely linked by covalent pairs. Figure 2 shows the diamond’s structure in 2-D and 3-D illustrations.

Figure 2. Diamond structure in 2-D and 3-D.

 

Another example is the silicon atom, of widespread use in electronic applications. It has a valence of four and obtains the eight electrons in its valence shell by sharing its electrons with four surrounding silicon atoms. Each occurrence of sharing signifies one covalent bond; thus, each silicon atom is bonded to four neighboring atoms by four covalent bonds.

 

The Metallic Bond

A metallic bond model is not as simple to construct as for the ionic and covalent cases. However, a simplified concept will be sufficient. In this kind of bond, the outer electrons are weakly held. If there are only a few valence electrons in an atom, they can be easily removed while the rest adhere firmly to the nucleus. This process forms a structure of positive ions and free electrons.

The positively charged atom cores consist of the nucleus and the remaining electrons. Because the valence electrons are free to move within the metallic structure, they form the commonly called electron cloud or gas. Positive ions and free electrons provide the attractive forces by which the metal atoms are held together.

Copper, for example, gives up its single valence electron, leaving behind a core consisting of the nucleus and inner electrons. It has a positive charge of one since one negatively charged electron is missing from the core.

 

Figure 3 is a schematic representation of free electrons – electron clouds – shared among a structure of positively charged copper atom cores (Z=29).

Figure 3. Metallic bond for copper.

 

The Crystal Structure

When atoms of a given type form a solid, they most frequently take up an orderly three-dimensional arrangement called a crystal.

All solids have the characteristics of defined size and shape. The definite shape and size require that the atoms or molecules of a solid be more or less in a fixed position and corporeity requires an efficient packing.

In its most stable form, the crystal, a solid has the additional property of regularity. The crystalline solids have a regularly repeated arrangement. Ideally, this order is so perfect that it is possible to specify each atom’s position in a crystal containing billions of atoms just by knowing a few atoms’ positions. The spacing between atoms in the crystal is the lattice constant.

We can identify a fundamental building block for each possible type of crystal, which can be repeated in space to generate as large a crystal as we please.

Germanium and silicon have a crystalline structure, with regular three-dimensional repetition of a unit cell looking like a tetrahedron with an atom at each vertex. Figure 4 shows a silicon crystal where each atom shares two electrons in a covalent bond with four of its neighbors.

Figure 4. 3-D representation of a silicon crystal.

 

A silicon piece may be a single crystal, or many individual crystals oriented randomly and joined together irregularly at their boundaries. The latter case is a polycrystalline form. Amorphous solids are those with randomly distributed atoms.

The best electronic devices employ single crystals, and one of the manufacturing problems is to obtain a very pure and regular crystal. Polycrystalline and amorphous materials also have applications in electronics.

 

About Interatomic Bonding in the Solid-State

The properties of the materials employed in electronics may be explained, on some occasions, by the type of interatomic bonding. For this reason, it is essential to have an understanding of the interatomic bonding types found in solids.

The primary atomic bonds in materials are ionic, covalent, and metallic. They are strong bonds, but, generally, metallic bonds are weaker than ionic and covalent bonds. 

In the ionic bonds, the transference of valence electrons from one atom type to another make electrically charged ions; forces are coulombic. NaCl or table salt is an example of an ionic bond.

In the covalent bond, there is a sharing of valence electrons between adjacent atoms. Diamond’s carbon atoms make covalent bonds with one another.

In metallic bonding, the valence electrons form an electron cloud or gas uniformly

diffused around the metal ion cores, acting as a kind of glue. The metallic bond is in charge of the cohesion of solid metals like copper, silver, and sodium.

The atoms that do not form ionic or covalent bonds attract themselves by weak electrostatic forces called Van der Waals forces – these are secondary bonds.

The crystalline solids have the atoms assembled in a specific order, forming a periodic array over large atomic distances. After solidification, the atoms locate themselves in repetitive three-dimensional patterns, where each atom bonds to its nearest-neighbor atoms. Under normal solidification conditions, metals, specific polymers, and many ceramic materials establish crystalline structures.

This long-range atomic order is absent in those materials that do not crystallize; they are noncrystalline or amorphous materials.