The Solid State

States of Matter of Class 11

The Solid State

Solids are rigid and have definite shapes. They also possess definite volume which is independent of the volume of the vessel.

Solids possess higher densities and are almost incompressible. All these characteristic properties are due to strong inter particle forces, smaller inter particle spares and restricted motion of the particles in the solid state.

The temperature at which solid changes into liquid state at normal pressure is called melting point of the substance.

Classification of Solids

Solids are broadly classified into two types: Crystalline solids and amorphous solids.

Distinction between Crystalline and Amorphous Solids

Sl. No.	Crystalline Solids	Amorphous Solids
1.	The internal arrangement of particles is regular	The internal arrangement of particles is not regular
2.	They have sharp melting points.	They do no have sharp melting points.
3.	Crystalline solids are regarded as true solid	Amorphous solids are regarded as super cooled liquids or pseudo solids
4.	Crystalline solids give a regular cut when cut with sharp edged knife.	Amorphous solids give irregular cut.
5.	Crystalline solids are anisotropic.	Amorphous solids are isotropic.

Classification of Solids based on different binding forces

Crystalline solids can be classified into different categories depending upon the type of constituent particles and the nature of attractive forces operating between them. Various categories are:

Atomic solids

Molecular solids
Ionic solids
Covalent solids
Metallic solids

Atomic Solids

In these solids the constituent particles are atoms. These closely packed atoms are held up by London dispersion forces. Some examples are crystals of noble gases. Such solids are very soft, possess very low melting points and poor conductors of heat and electricity.

Molecular Solids

In these solids, the constituent particles which pack up together are molecules of the substance. These molecules may be non – polar (dipole moment = 0) such as CO₂,I₂,CCl₄ etc. or they may be polar (dipole moment > 0) like H₂O, HCl, HF, etc.

In case of non – polar molecules, the attractive forces operating between the molecules are Vander Waal forces (also called dispersion forces). The example of such solids are : dry ice (Solid CO₂, iodine (crystals).
In case of polar molecules, the attractive forces operating between the molecules in solid state are dipole – dipole forces. The examples of such solids are : solid S₂O, solid HCl. However, in some solids with polar molecules, the interparticle forces are hydrogen bonds. The examples of such solids are ice (H₂O) solid hydrogen fluoride (HF); solid ammonia (NH₃), etc.

Characteristics of Molecular Solids

Some of the general characteristics of molecular solids are :

They are generally soft.

Their melting points are low to moderately high. The melting points of solids with non – polar molecules are relatively low whereas solids with polar molecules have moderately high melting points.

They are generally bad conductors of heat and electricity.

They have generally low density.

Ionic Solids

In ionic solids, the constituent particles are ions of opposite charges. Each ion is surrounded by a definite number of ions of opposite charge. The number of ions that surround a particular ion of opposite charge its called co – ordination number of the ion. For example, in sodium chloride crystal each sodium ion (Na⁺) is surrounded by six chloride ions. Hence coordination number of Na⁺ is 6. At the same time each chloride ion is surrounded by six Na⁺ ions. Therefore the co – ordination number of Cl^- ion is also 6. However, in calcium fluoride crystal each Ca²⁺ ion is surrounded by eight fluoride (F^-) ions and each F^- ion is surrounded by four Ca²⁺ ions. Thus, in CaF₂ crystal co – ordination numbers of Ca²⁺ and F^- ions are respectively 4 and 8. The interparticle forces in ionic solids are ionic bonds operating between the ions of opposite charges some examples of ionic solids are : sodium chloride (NaCl) ; ceasium chloride (CsCl), zinc sulphide (ZnS), calcium fluoride (CaF₂), etc.

Characteristics of Ionic Solids

Some common characteristics of ionic solids are as follows:

They are hard, brittle and have low volatility.

They have high melting points.
They are poor conductors of electricity in solid state, however they become good conductors of electricity in molten state or in dissolved state.
They are generally soluble in polar solvents like water.

Covalent Solids

In these types of solids the constituent particles are atoms of same or different elements connected to each other by covalent bond network. For example, in diamond only carbon atoms constitute the covalent network while carborundum covalent bond network is constituted by silicon and carbon atoms. Obviously, the interparticle forces operating in these solids are covalent bonds. These solids are also called network solids because the covalent bonds extend in three dimensions forming a giant interlocking structure. Some examples of covalent solids are :

Diamond, silicon carbide, aluminium nitrite etc.

Characteristics of Covalent Solids

Some common characteristics of covalent solids are :

They are very hard. Diamond is the hardest naturally occurring substance.
They have very high melting points.
They are poor conductors of heat and electricity.
They have high heats of fusion.

Metallic Solids

In these type of solids, the constituent particles are metal atoms. The interparticle forces in these solids are metallic bonds. In the metallic crystals the metal atoms occupy the fixed positions but their valence electrons are mobile. The close packed assembly of metal kernels (part of metal atom without valence electrons) remain immersed in the sea of mobile valence electrons. The attractive force between the kernels and mobile valence electrons is termed as metallic bond.

Characteristics of Metallic Solids

The common characteristics of metallic solids are as follows:

They generally range from soft to very hard.
They are malleable and ductile.
They are good conductors of heat and electricity.
They possess bright lustre.
They have high melting and boiling points.
They have moderate heats of fusion.

The summary of classification of solids on the basis of interparticle forces is given in

Classification of Solids on the Basis of Binding Forces

Crystal Classification	Unit Particles	Binding Forces	Properties	Examples
Atomic	Atoms	London dispersion forces	Soft, very low melting, poor thermal and electrical conductors	Noble gases
Molecular	Polar or non – polar molecules	Vander Waal’s forces (London dispersion, dipole – dipole forces hydrogen bonds)	Fairly soft, low to moderately high melting points, poor thermal and electrical conductors	Dry ice (solidCO₂), methane (CH₄)
Ionic	Positive and negative ions	Ionic bonds	Hard and brittle, high melting points, high heats of fusion, poor thermal and electrical conductors	NaCl, ZnS
Covalent	Atoms that are connected in covalent bond network	Covalent bonds	Very hard, very high melting points, poor thermal and electrical conductors	Diamond, quartz, silicon
Metallic	Cations in electron cloud	Metallic bonds	Soft to very hard, low to very high melting points, excellent thermal and electrical conductors, malleable and ductile	All metallic elements, for example, Cu, Fe, Zn

Intermolecular forces or Vander Waal’s forces

Intermolecular forces or vander Waals’ forces originate from the following three types of interactions.

Dipole – Dipole interactions: In case of polar molecules, the vander waals’ forces are mainly due to electrical interaction between oppositively charged ends of molecules (Fig. 1. a) called dipole – dipole interactions. For example, gases such as NH_{3 ,}SO_2,HCl,HF etc.have permanent dipole moments as a result of which there is appreciable dipole – dipole interactions between the molecules of these gases. The magnitude of this type of interaction depends upon the dipole moment of the molecule concerned. Evidently, greater the dipole moment, stronger is the dipole – dipole interactions. Because of these attractive forces, these gases can be easily liquefied.

The Solid State

Dipole – Induced dipole Interactions: A polar molecule may sometimes polarize a non – polar molecule which lies in its vicinity and thus induces polarity in that molecule just as a magnet induces magnetic polarity in a neutral piece of iron lying close by. The induced dipole then interacts with the dipole moment of the first molecule and thereby the two molecules are attracted together (Fig. 1. b). The magnitude of this interaction, evidently depends upon the magnitude of the dipole moment of the polar molecule and the polarizability of the non – polar molecule. The solubility of inert gases in H₂O increases from He to Rn due to a corresponding increase in magnitude of the dipole – induced dipole interactions as the polarizability of the noble gas increases with increase in size from He to Rn.
Momentary dipole – induced dipole interactions: The electrons of neutral molecules keep on oscillating w.r.t. the nuclei of atoms. As a result, at a given instant, one side of the molecule may have a slight excess of electrons relative to the opposite side. Thus a non – polar molecule may become momentarily self – polarized. This polarized molecule may induce a dipole moment in the neighbouring molecule. These two induced dipoles then attract each other (Fig. 1. c). These momentary dipole – induced dipole attractions are also called London forces or dispersive forces. The magnitude of these forces depends upon the following:

(i) Size or molecular mass: The melting points and boiling points of non – polar molecules increase as the size or molecular mass of the molecule increases. For example, the m.p. and b.p. of alkanes, halogens, noble gases etc. increase as the molecular mass of the molecule increases.

(ii) Geometry / Shape : For example, isomer n – pentane has higher boiling point than neo – pentane because the former is zig – zag chain with larger sites of contact and hence large intermolecular forces whereas the latter is nearly spherical and hence has less contact and weaker forces.