In nature one comes across several types of solids. Many solids are aggregates of atoms. The arrangement of atoms in any solid material is determined by the character, strength and directionality of the chemical binding forces, cohesive forces or chemical bonds. We call these binding forces as atomic interaction forces. The atoms, molecules or ions in a solid state are more closely packed than in the gaseous and liquid states and are held together by strong mutual forces of attraction and repulsion. One can describe the atomic arrangement in elements and compounds on the basis of this.
Types of Bonds
Seitz in 1940 classified solids into five types according to the bonding of atoms, which has becomes a generally adapted classified
Table below Classification of solids according to the bonding of atoms
Types of bond
Binding energy (eV/atom)
Covalent, atomic or homopolar bonds
Electron shared between two atoms
Carbon (diamond) Ge, Si, SiC, BN etc
Ionic or electrostatic bonds
Electron transfer and Coulomb interaction between cations and anions
Freely moving electrons in an array of positive ions
All metals and alloys
4. Molecular (Vander Waals)
Molecules between pairs
Weak attractive forces due to dipole-dipole interaction
5. Hydrogen bond
Hydrogen atom attracted between two other atoms
Electrostatic bond of H-atom with an electronegative atom
Ice, organic compounds, biological materials
The atoms of different types come closer and join together during chemical reaction and usually these bonds are referred to as chemical bonds.
FORCES BETWEEN ATOMS: MECHANISM OF BOND FORMATION AND BOND ENERGY
We have already stated that the interatomic forces exist amongst the atoms of a crystal and are responsible for holding atoms together to form solid structure. We have also said that the, forces between atoms can be of two kinds: (i) attractive forces which keep the atoms together, and (ii) repulsive forces which come into play when the solid is compressed. The force of attraction between atoms brines them closer together until the individual election clouds begin to overlap and a strong repulsive force arises to comply with Pauli’s exclusion principle. The attractive forces decrease the potential energy of the system and repulsive forces increase it Obviously, the net energy of the system is equal to the algebraic mm of these two energies. When the attractive force and the repulsive force between any two atoms are equal, the two atoms should be in a stable situation with a minimum potential energy. The forces between two atoms or ions in a solid as a function of their distance of separation r are shown in Figure below.
The separation distance (r) between the centre to centre of two bonding atoms which gives the stable bond is called bond length. Obviously, greater the force of attraction between the two bonding atoms, smaller will be the bond length (r). We must note that the primary bonds are stronger than the secondary bonds. One can use the bond length for finding the diameter of ions or atoms as follows:
r = rc + ra, where rc → radii of a cation and ra → radii of anion
Perhaps ionic or hetropolar bonding which is formed by the actual transfer of electrons from one atom to the other so that each atom acquires a stable configuration similar to the nearest inert gas atoms, is the simplest type of chemical bonding to visualize since it is almost totally electrostatic in nature. Ionic bonding occurs between electropositive elements (metals, i.e., those elements on the left side of the periodic table) and the electronegative elements (non-metals, i.e. on the right hand of the periodic table). These bonds are formed mainly in inorganic compounds, e.g., sodium chloride (common salt NaCl), MgO, CuO, CrO2, and MOF2. In MgO the ions are doubly ionized leading to a stronger interatomic bond and hence a higher melting point (), compared to 800°C for NaCl.
An ionic bond is truly speaking the attractive force existing between a positive and negative ion when they are brought into close proximity. These ions, of course, are formed when the atoms involved lose or gain electrons in order to stabilise their outer shell electron configuration. The cohesive energy of ionic crystal is quite large, it is of the order of 5 to 10 eV.
A good example of ionic crystals is the crystal of sodium chloride (NaCl). The electronic configurations of white soft metal Na and Cl atoms are as follows :
Na : 1s2 2s2 2p6 3 s1 [K (2) L (8) M(1)]
Cl : 1s2 2s2 2p6 3s2 [K (2) L (8) M(7)]
Na atom has a low ionisation energy and hence easily loses an electron; where as C1 atom has a high electron affinity and strongly tends to acquire an electron. The reaction can be represented by
Na + Cl → Na+ + Cl- → NaCl
Since chlorine exists as molecules, we can write the chemical reaction as
2Na + Cl2 → 2Na+ + 2Cl- → 2NaCl
After the transfer of an electron from 3s orbital of Na to 3p orbital of C1, the electronic configurations become Na+ : 1s2 2s2 2p6 (same as Ne)
Cl- : 1s2 2s2 2p6 3s2 3p6 (same as Ar)
This type of bonding is formed by an equal sharing of electrons between two neighbouring atoms each having incomplete outermost shells. The atoms do so in order to acquire a stable electronic configuration in accordance with the octet rule. Unlike ionic bonding the atoms participating in the covalent bond have such electronic configurations that they cannot complete their octets transfer of electrons from one atom to the other. Obviously, there is no charge associated with any atom of the crystal, The majority of solids incorporating covalent bonds are also bound by either ionic or Vander Waals bonds.
Characteristics of Covalent Compounds
(i) Covalent compounds are mostly gases and liquids
(ii) They are usually electric insulators.
(iii) They are directional in nature.
(iv) They are insoluble in polar solvents like water but are soluble in non-polar solvents, e.g., benzene, chloroform, alcohol, paraffin sets,
(v) Covalent compounds are homopolar, i.e. the valence electrons are bound to individual or pairs of atoms and electrons cannot move freely through the material as in the case metallic bonds.
(vi) Covalent compounds are soft, rubbery elastomers, and form a variety of structural materials usually termed as plastics. The melting and boiling points of these compounds are low.
(viii) In the exceptional case of diamond covalent bonding is very strong due to the very large cohesive energy. Usually the binding energy ranges from 1 to 5 eV. The strong covalent bonding makes diamond very hard and with a high melting point.
Metallic bond is formed by the partial sharing of valence electrons by the neighbouring atoms. They are formed by the elements of all subgroups A and I-II, subgroups B. Unlike the case of covalent bond, the sharing in metallic bond is not localized. Hence metallic bond nay considered as delocalized or unsaturated covalent bond. Metallic bonds are electropositive. When interacting with elements of other groups, atoms in a metallic crystal can easily give off their valence electrons and change into positive ions.
Comparison between ionic, covalent and metallic bonds
This type of bonding exists due to electrostatic force of attraction between positive and negative ions of different elements.
This type of bonding exists due to the electrostatic force of attraction between atoms which share the electron pairs to form a covalent bond.
This type of bonding exists due to electrostatic force of attraction between electron clouds and positive ions of same or different metals.
This type of bonding is formed between two different elements. One of the atoms loses its valence electron and other accept it. The ions so formed attract each other and an ionic bond is formed.
This type of bonding is formed due to sharing of electron pairs between the atoms of same or different elements.
When the valence electrons detach themselves from their parent atoms and form a common pool, metallic bond is formed. The forces which bind the electron cloud and positive ions of the metal forms the metallic bond.
Due to their crystalline structure, ionic solids have high hardness.
Except diamond, silicon, and carbide, etc, the covalent solids have low hardness.
The metallic solids are soft in nature and have crystalline structure.
The ionic solids have very low thermal and electrical conductivities.
The covalent solids have low thermal and electrical conductivities.
The metallic solids have very high thermal and electrical conductivities.
The ionic solids have high melting and boiling temperatures.
The covalent solids have lower melting and boiling temperatures in comparison to ionic solids.
In comparison to ionic solids, metallic solids have slightly lower and boiling temperatures.
The ionic solids are not malleable and ductile.
The covalent solids are not soluble in toluene and benzene etc.
The metallic solids are not malleable and ductile.
The ionic solids are soluble in water.
The covalent solids are soluble in toluene and benzene etc.
The metallic solids are neither soluble in water nor in benzene, etc.
The ionic compounds exist in the solid form only.
The covalent compounds exist in solid, liquid, and gaseous form.
Except mercury, metallic compounds exist in solids form.
This is a special type dipole bond, but it is considerably stronger. It is a special case of inter-molecular attraction produced between certain covalently bonded hydrogen atoms and one pairs of electrons of another atom. When a hydrogen atom is covalently bonded to a relatively large atom such as nitrogen, oxygen or fluorine, a powerful permanent dipole is set up. There is a positive field adjacent to the hydrogen and a negative field around the electron pair. One finds that this force is very important in plastics and in biological molecules such as DNA. A good example of hydrogen bond is water molecule.
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