Bonds in Metals: 4 Types | Metallurgy

The following points highlight the four main types of bonds in metals. The types are: 1. Ionic Bond 2. Covalent Bond 3. Van Der Waals Bond 4. Metallic Bond.

Atoms are bonded with each other in solids by four important mechanisms. In three of these, bonding is made when the atoms fill their outer ‘s’ and ‘p’ orbitals.

Type # 1. Ionic Bond:

When a solid has more than one type of atoms, then one atom donates its valence electrons to a different atom, filling the outer energy shell of the second atom. By this transfer of electrons, though both atoms now have filled (or empty) outer energy levels (stable state), but both have acquired an electrical charge and thus, behave as ions.

The first atom which donates the electrons is left with a net positive charge, and thus, is a cation, while the second atom which accepts the electrons acquires a net negative charge, and thus, is an anion. The oppositively charged ions thus have electrostatic forces between them and are attracted to one another and thus produce the ionic bond. Fig. 1.1 illustrates the classic example of sodium chloride. The attraction between sodium and chloride ions produces the solid sodium chloride.

The electrical conductivity of ionic bonded solids is poor. The electrical charge can be transferred by the movement of ions, which because of their size do not move as easily as electrons (in metallic bond).

Moreover as an ion tries to move, it faces large repulsive forces due to coming closer to similarly charged ions. Aqueous solutions or fused ionic solids are good conductors of electricity, because then, the ions can move easily and transfer the electrical charge.

An ionic crystal consists of an aggregate of a large number of ions packed together forming a three dimensional lattice in which each ion is surrounded by ions of opposite charge. Fig. 1.1 (c) illustrates that each Na⁺ ion is surrounded by 6 Cl^– ions, i.e., each ion attracts all its neighbouring ions of opposite charge.

The ionic bond is non-directional (The actual number of surrounding neighbours of an ion depends on first by geometric factors, and second by the necessity of preserving electrical neutrality in the solid), and is a strong bond.

When subjected to stresses, ionic crystals tend to cleave (break) along certain planes of atoms rather than to deform in a ductile manner as metals do.

Type # 2. Covalent Bond:

In covalently bonded solids, sharing of electrons among two or more atoms takes place. Actually, covalent bonding requires that electrons be shared between atoms in such a way that each atom gets its outer sp orbital filled.

Let us take an example of silicon. A silicon atom has four valence electrons, Fig. 1.3 (a). It acquires its octet (eight electrons) in its outer energy shell by sharing the electrons with four surrounding silicon atoms, Fig. 1.3 (b). Each, sharing represents one covalent bond, and thus, each silicon atom is bonded to four neighbouring atoms by four covalent bonds.

For the formation of covalent bonds, the silicon atoms have to be arranged in space so that the bonds have a fixed directional relationship with one another. Thus, the arrangement of atoms results in a tetrahedron with angles of 109.5° between the covalent bonds, Fig. 1.3 (c). To complete the octet of electrons needed for atomic stability, the electrons must be shared with 8-N (8 minus N) neighbouring atoms, where N is the number of valence electrons in the given element.

Non-metallic elements like nitrogen, carbon, oxygen, fluorine and chlorine have covalent bonds. Si, Se, Ge, As have partly covalent and partly metallic bonds. Transition metals too have partly covalent bonding.

The sharing of electrons (or overlapping of orbitals) makes the covalent bonds to be very strong. Diamond, Fig. 1.4 is the hardest natural material. But covalently bonded solids have typically poor ductility (a glass shatters when dropped), and poor electrical and thermal conductivity (a brick is a good insulator). For the electrons to move and carry a current requires the covalent bond to be broken, Fig. 1.5 (a), which requires high temperatures and application of high voltage.

Type # 3. Van Der Waals Bond:

It is a secondary bonding as it joins (bonds) molecules, or group of atoms by weak electrostatic attraction, but the atoms within the molecules, or group of atoms are bonded by strong covalent or ionic bonds. For example, heating water to its boiling point (≈ 100°C, a very low temperature) breaks the van der Waals bonds in between the water molecules, and changes water to steam. But heating too much higher temperatures is needed to break the covalent bonds joining oxygen and hydrogen atoms of each water molecule.

Many plastics, ceramics, water and other molecules are permanently polarised, i.e., some parts of a molecule are positively charged while other parts are negatively charged (due to more possibilities of electrons staying there in the latter parts and conversely for the former).

Fig. 1.6 illustrates that this causes weak electrostatic attractions between the positively charged parts of one molecule and the negatively charged parts of other water molecules. Weak bonds bind the two molecules. This weak van der Waal bonding is also called hydrogen bonding if hydrogen atom (such as in between water molecules) represents one of the polarised parts.

The nature of these weak bonds can change dramatically the properties of materials and is often taken advantage of. For example, Polyvinyl chloride (PVC) has chain like long molecules.

The bonding within each chain is strong covalent, but individual chains are bonded to one another by weak van der Waal bonds. As these bonds can be easily broken to make the chains to slide past one another, and thus polyvinyl chloride can be deformed, i.e., becomes plastic.

Type # 4. Metallic Bond:

Metals are aggregate of atoms, usually arranged in a crystal structure. The properties of the metal depend not only on the nature of constituent atoms, but also as to how are they bonded.

In a metal (or an alloy) which has a low electronegativity, each of the atoms contributes its valence electrons to the formation of an electron cloud surrounding the ions. Fig. 1.7 (b) illustrates the nature of the metallic bond having ions embedded in a cloud of electrons.

The electrons in this cloud are free to move to any part and in virtually any direction. Such a cloud of valence electrons are called free electrons. For example, aluminium atom contributes its three valence electrons leaving behind a core consisting of the nucleus and inner electrons.

As three negatively charged electrons are missing from the atom, the core has a positive charge of three. The valence electrons join the electron sea and become associated with several cores of atoms. The positively charged-atom-cores are held together by mutual attraction to the nearby electrons, Fig. 1.7 (b), thus, producing a strong metallic bond.

Metallic bond is sometimes considered to have characteristics similar to ionic as well as covalent bonds. It resembles ionic bond if one imagines that the negative electrons are holding positive ions together, whereas it resembles the covalent bond as the electrons are shared by adjacent atoms.

But metallic bond can exist among a large aggregate of atoms, while a covalent bond can occur between as few as two atoms. Also, metallic bonds are non-directional as each valence electron is not localised between only two ion cores (as in covalent bonding), but are more or less free to travel throughout the solid metal.

In general, fewer the valence electrons an atom has and more loosely they are held, the more is the metallic bonding. For example, metals like sodium, potassium, copper, silver and gold have high electrical and thermal conductivity, because the few valence electrons present are highly mobile.

As the number of valence electrons and the tightness with which they are held to the nucleus, increases, the valence electrons become more localised (i.e., the mobility throughout the solid becomes less).

This increases the covalent nature of the bonding. In elements of the fourth group of periodic table, diamond (non-metal) exhibits almost pure covalent bonding; Silicon atoms covalently share electrons, but a few electrons are able to leave the covalent location between adjacent atoms to permit limited conductivity.

This occurs more extensively in tin and germanium, and thus acts as a base for semi-conductivity. These metals have more metallic behaviour. Tin actually occurs in two modifications, one mostly covalent and the other mostly metallic. Lead is mostly metallic.

The transition metals (atoms of which have incomplete ‘d’ shell such as iron, nickel and tungsten) have large fraction of covalent bonding involv­ing hybridized inner shell electron orbitals, which explains in part their high melting points. In fact, it is seldom to come across pure metallic bonds in a metallic material, but in some combination of the four major types of bonding. Fig 1.8 illustrates a tetrahedron with these pure bonds at the apexes of it and the combination of bonding for usual metallic materials.

For example, iron is bonded by a combination of metallic and covalent bonding, which prevents atoms from packing as efficiently as could be otherwise. Compounds formed between two or more metals, called intermetallic compounds may have a combination of metallic and ionic bonding, especially when there is a large difference in electronegativity between elements.

As the electronegativity difference between the atoms increases, the bonding becomes more ionic. For example, both aluminum and vanadium have electronegativity of 1.5, the compound, Al₃V has primarily metallic bonds. But lithium has electronegativity of 1.0, and thus, the compound AlLi has a combination of metallic and ionic bonding.