Solid State, Class 12, Chemistry Detailed Chapter Notes PDF Download

Amorphous and Crystalline Solids and Classification - The Solid State, Class 12, Chemistry

Solid Liquid Gas

 (1) Have definite volume Have definite volume Have indefinite volume

(2) Their shape is fixed Their shape is indefinite volume Their shape is indefinite

(3) They have strongest Their force of attraction is They have very weak force of intermolecular attraction. intermediate between solids attraction . Hence molecules are

Hence their shape as well and gases . So shape is very - very loosely connected

as volume is fixed variable but volume is fixed. with each other and hence have neither shape nor volume fixed.

The outstanding macroscopic properties of gases are compressibility and fluidity.

In contrast, the most noticeable macroscopic features of crystalline solids are rigidity, incompressibility and characteristic geometry. We shall find that the explanation of these macroscopic properties in terms of the atomic theory involves the idea of lattice: a permanent ordered arrangement of atoms held together by forces of considerable magnitude.

Thus the extremes of molecular behaviour occur in gases and solids. In the former we have molecular chaos and vanishing intermolecular forces, and in the latter we have an ordered arrangement in which the interatomic forces are large.

Types of Solids

a) Crystalline solids: In a single crystal the regularity of arrangement of the pattern extends throughout the solid and all points are completely equivalent.

b) Amorphous solids: An amorphous solid differs from a crystalline substance in being without any shape of its own and has a completely random particle arrangements, i.e. no regular arrangement. Example: Glass, Plastic

Classification of crystalline

Molecular solids : Those solids which consists of small molecules are called molecular solids.

(1) Non - polar molecular solids :

(electronegativity sequence (F > O > N ≈ Cl > Br > S > C H)

The solids which have zero dipole moment are called non-polar molecular solids.

The molecules are held together by weak Vander waal's forces . Hence they are either gas or liquids at room temperature.

They are poor conductor of electricity due to their non-polar nature.

(2) Polar molecular solids :

Those solids which have non-zero dipole moments are called polar molecular solids.

Polar molecular solids have dipole - dipole interaction, which is slightly stronger than vander waal's force and hence they have larger melting and boiling points than the non - polar molecular solids .

Example : Solid CO₂ , Solid NH₃

They are generally liquids or gases at room temperature.

(3) Hydrogen bonded molecular solids :

Those molecular solids which are bonded with each other by hydrogen

bonds are called hydrogen bonded molecular solids.

Example : Ice

They are non-conductor of electricity

Generally they are liquid at room temperature or soft solids .

IONIC SOLIDS : All those solids whose constituent particles are ions are called ionic solids,

Example : NaCl, CsBr, AgBr, CsCl

 METALLIC SOLIDS :

• All those solids which are bonded by metallic bonds are called metallic solids.
• Inner core electrons are immobile
• Metallic solids show a great electrical conductivity due to availability of large number of free electrons whose movements constitute electric current.
•  Metallic solids are malleable and ductile.
•  Metallic solids have lustre.
•  Metallic solids have good thermal conductivity.

COVALENT SOLIDS OR NETWORKED SOLIDS :

Whenever electric field is applied, electrons move between the layers.

Graphite is a good conductor of electricity due to availability of free electrons.

Networked solids are hard and brittle.

Carbon - carbon bond has got partial double bond character in graphite.

Two layers in graphite are attached with each other by weak vander waal force.

Graphite can be used as a lubricant at high temperatures .

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Crystal Lattice and Unit cells - The Solid State, Class 12, Chemistry

What is a crystal:

A crystalline solid consist of a large number of small units, called crystals, each of which possesses a definite geometric shape bounded by plane faces. The crystals of a given substance produced under a definite set of conditions are always of the same shape.

Unit Cells

In this topic we would be studying certain properties of a solid which depend only on the constituents of the solid and the pattern of arrangement of these constituents. To study these properties, it would be convenient to take up as small amount of the solid as possible because this would ensure that we deal with only the minimum number of atoms or ions. This smallest amount of the solid whose properties resemble the properties of the entire solid irrespective of the amount taken is called a unit cell. It is the smallest repeating unit of the solid. Any amount of the solid can be constructed by simply putting as many unit cells as required.

Bravais Lattices

Bravais (1848) showed from geometrical considerations that there are only seven shapes in which unit cells can exist. These are: (i) Cubic (ii) Orthorhombic (iii) Rhombohedral (iv) Hexagonal (v) Tetragonal (vi) Monoclinic and (vii) Triclinic. Moreover he also showed that there are basically four types of unit cells depending on the manner in which they are arranged in a given shape. These are : Primitive, Body Centered, Face Centered and End Centered. He also went on to postulate that out of the possible twenty eight unit cells (i.e. seven shapes ´ four types in each shape = 28 possible unit cells), only fourteen actually would exist. These he postulated based only on symmetry considerations. These fourteen unit cells that actually exist are called Bravais Lattices.

Primitive Cubic Unit Cell

In a primitive cubic unit cell the same type of atoms are present at all the corners of the unit cell and are not present anywhere else. It can be seen that each atom at the corner of the unit cell is shared by eight unit cells (four on one layer, as shown, and four on top of these). Therefore, the volume occupied by a sphere in a unit cell is just one-eighth of its total volume. Since there are eight such spheres, the total volume occupied by the spheres is one full volume of a sphere. Therefore, a primitive cubic unit cell has effectively one atom. A primitive cubic unit cell is created in the manner as shown in the figure. Four atoms are present in such a way that the adjacent atoms touch each other. Therefore, if the length of the unit cell is `a', then it is equal to 2r, where r is the radius of the sphere. Four more spheres are placed on top of these such that they eclipse these spheres. The packing efficiency of a unit cell can be understood by calculating the packing fraction. It is defined as ratio of the volume occupied by the spheres in a unit cell to the volume of the unit cell and void fraction is given as 1 Packing fraction. Therefore, PF = . (This implies that 52 % of the volume of a unit cell is occupied by spheres). VF » 0.48

Figure 3(b)

Body Centered Cubic Unit Cell

A Body Centred unit cell is a unit cell in which the same atoms are present at all the corners and also at the center of the unit cell and are not present anywhere else.This unit cell is created by placing four atoms which are not touching each other. Then we place an atom on top of these four. Again, four spheres eclipsing the first layer are placed on top of this. The effective number of atoms in a Body Centered Cubic Unit Cell is 2 (One from all the corners and one at the center of the unit cell). Moreover, since in BCC the body centered atom touches the top four and the bottom four atoms, the length of the body diagonal () is equal to 4r. The packing fraction in this case is = VF » 0.32

figure 4(b)

Face Centered Cubic Unit Cell

In a fcc unit cell, the same atoms are present at all the corners of the cube and are also present at the centre of each square face and are not present anywhere else. The effective number of atoms in fcc is 4 (one from all the corners, 3 from all the face centers since each face centered atom is shared by two cubes). Since, here each face centered atom touches the four corner atoms, the face diagonal of the cube () is equal to 4r.

VF » 0.26

Figure 5(b)

Hexagonal Primitive Unit Cell

Each corner atom would be common to 6 other unit cells, therefore their contribution to one unit cell would be 1/6. Therefore, the total number of atoms present per unit cell effectively is 6.

Figure 6(b) ABCD is the base of hexagonal unit cell AD=AB=a. The sphere in the next layer has its centre F vertically above E and it touches the three spheres whose centres are A,B and D.

Figure 6(a)

Hence FE = =

The height of unit cell (h) =

The area of the base is equal to the area of six equilateral triangles, =  . The volume of the unit cell = .

Therefore

PF = = 0.74 VF » 0.26

Density of crystal lattice

The density of crystal lattice is same as the density of the unit cell which is calculated as

r = =

r =

Seven crystal systems : The seven crystal systems are given below.

Illustration 1: Lithium borohydride crystallizes in an orthorhombic system with 4 molecules per unit cell. The unit cell dimensions are a = 6.8Å,

b = 4.4Å and C = 7.2Å. If the molar mass is 21.76g. Calculate density of crystal.

Solution: Since

Density, r =

Here n = 4, Av. No = 6.023 ´ 10²³, &

Volume = V = a x b x´ c

= 6.8 x 10⁸ x 4.4 x 10⁸ x 7.2 x 10⁸ cm³ = 2.154 ´ 10²² cm³

density, r = = 0.6708 gm/cm³

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Packing Efficiency in Cubic-closed Packed Structures - The Solid State, CBSE, Class 12, Chemistry

Close Packing of Spheres

Atoms are space filling entities and structures can be described as resulting from the packing of spheres. The most efficient, called Closest Packing, can be achieved in two ways, one of which is called Hexagonal Close Packing (Hexagonal Primitive) and the other, Cubic Close Packing (ccp or fcc). Hexagonal close packing can be built up as follows:

Place a sphere on a flat surface. Surround it with six equal spheres as close as possible in the same plane. Looking down on the plane, the projection is as shown in Fig. 7 (a). Let us call this layer as the A layer. Now form over the first layer a second layer of equally bunched spheres, Fig 7 (b), so as to nestle into the voids. It will be clearly seen that once a sphere is placed over a void, it blocks the void which are adjacent to that void. Let us call this layer as the B layer. Now, it can be clearly seen that there are two types of voids created by the B layer of spheres. If a sphere is placed on the x - type of voids, it would resemble the A layer of spheres in the sense that it eclipses the spheres of the A layer. This arrangement (i.e., ABAB……) is called the hexagonal close packing (hcp) or hexagonal primitive. On the other hand, if the spheres were to be placed on the y - type of voids, it would neither eclipse the A layer nor the B layer of spheres. This would clearly be a unique layer. Let us call this layer as the C layer (as seen in figure 7(c). This arrangement (i.e, ABCABC...) is called the Cubic Close Packing (fcc).

Exercise 1: It can be seen now that both fcc and Hexagonal Primitive Structure have the same packing fraction. Moreover this is also the highest packing fraction of all the possible unit cells with one type of atom with empty voids. Can you explain this?

Octahedral and Tetrahedral Voids

The close packing system involves two kinds of voids - tetrahedral voids and octahedral voids. The former has four spheres adjacent to it while the latter has six spheres adjacent to it. These voids are only found in either fcc or Hexagonal Primitive unit cells. Let us first consider a fcc unit cell (as shown in the figure 8(a).). Not all the atoms of the unit cell are shown (for convenience). Let us assume that there is an atom (different from the one that forms the fcc) at the center of an edge.

Let it be big enough to touch one of the corner atoms of the fcc. In that case, it can be easily understood that it would also touch six other atoms (as shown) at the same distance. Such voids in an fcc unit cell in which if we place an atom it would be in contact with six spheres at equal distance (in the form of an octahedron) are called octahedral voids (as seen in fig. 8(a).

On calculation, it can be found out that a fcc unit cell has four octahedral voids effectively. The number of effective octahedral voids in a unit cell is equal to the effective number of atoms in that unit cell.

Let us again consider a fcc unit cell. If we assume that one of its corners is an origin, we can locate a point having coordinates (¼,¼,¼). If we place an atom (different from the ones that form the fcc) at this point and if it is big enough to touch the corner atom, then it would also touch three other atoms (as shown in the fig. 8(b) which are at the face centers of all those faces which meet at that corner. Moreover, it would touch all these atoms at the corners of a regular tetrahedron. Such voids are called tetrahedral voids (as seen in fig. 8(b).

Since there are eight corners, there are eight tetrahedral voids in a fcc unit cell. We can see the tetrahedral voids in another way. Let us assume that eight cubes of the same size make a bigger cube as shown in the fig. 8(c). Then the centers of these eight small cubes would behave as tetrahedral voids for the bigger cube (if it were face centered). The number of tetrahedral voids is double the number of octahedral voids. Therefore, the number of tetrahedral voids in hcp is 12.

Radius Ratio Rules

The structure of many ionic solids can be accounted for by considering the relative sizes of the positive and negative ions, and their relative numbers. Simple geometric calculations allow us to workout, as to how many ions of a given size can be in contact with a smaller ion. Thus, we can predict the coordination number from the relative sizes of the ions.

Co-ordination Number 3

The adjacent fig. 9(a) shows the smaller positive ion of radius r⁺ is in contact with three larger negative ions of radii r^-. It can be seen that AB = BC = AC = 2r^-, BD = r^- + r⁺. Further, the angle ABC is 60^o, and the angle DBE is 30^o. By trigonometry Cos 30^o = (BE / BD). BD = (BE / Cos 30^o), r⁺ + r^- = r^- / Cos 30^o = (r^- / 0.866) = r^- ´ 1.155, r⁺ = (1.155 r^-) - r^- = 0.155 r^-, Hence (r⁺ / r^-) = 0.155.

Co-ordination Number 4 (Tetrahedron)

Angle ABC is the tetrahedral angle of 109^o 28' and hence the angle ABD is half of this, that is 54^o44'. In the triangle ABD, Sin ABD = 0.8164 = AD / AB = . Taking reciprocals, , rearranging, we get, .

Co-ordination Number 6 (Octahedron) or 4 ( Square Planar)

A cross section through an octahedral site is shown in the adjacent figure and the smaller positive ion (of radius r⁺) touches six larger negative ions (of radius r^-). (Note that only four negative ions are shown in this section and one is above and one below the plane of the paper). It is obvious that AB = r⁺ + r^- and BD = r^-. The angle ABC is 45^o in the triangle ABD. cos ABD = 0.7071 = (BD / AB) = r. Rearranging, we get, r⁺/r= 0.414

Exercise 2: The radius of a calcium ion is 94 pm and of an oxide ion is 146 pm. Predict the crystal structure of calcium oxide.

Classification of Ionic Structures

In the following structures, a black circle would denote a cation and a white circle would denote an anion.In any solid of the type A_xB_y, the ratio of the coordination number of A to that of B would be y : x.

Rock Salt Structure

Cl^- is forming a fcc unit cell in which Na⁺ is in the octahedral voids. The coordination number of Na⁺ is 6 and therefore that of Cl^- would also be 6. Moreover, there are 4 Na⁺ ions and 4 Cl^- ions per unit cell. The formula is Na₄Cl₄ i.e,NaCl. The other substances having this kind of a structure are halides of all alkali metals except cesium halides and oxides of all alkaline earth metals except beryllium oxide.

Figure 10 Unit cell structure of NaCl

Zinc Blende Structure

Sulphide ions are face centered and Zinc is present in alternate tetrahedral voids. Formula is Zn₄S₄, i.e, ZnS. Coordination number of Zn is 4 and that of sulphide is also 4. Other substance that exists in this kind of a structure is BeO.

Fluorite Structures

Calcium ions are face centered and fluoride ions are present in all the tetrahedral voids. There are four calcium ions and eight fluoride ions per unit cell. Therefore the formula is Ca₄F₈, (i.e, CaF₂). The coordination number of fluoride ions is four (tetrahedral voids) and thus the coordination number of calcium ions is eight. Other substances which exist in this kind of structure are UO₂, and ThO₂.

Anti-Fluorite Structure

Oxide ions are face centered and lithium ions are present in all the tetrahedral voids. There are four oxide ions and eight lithium ions per unit cell. As it can be seen, this unit cell is just the reverse of Fluorite structure, in the sense that, the positions of cations and anions is interchanged. Other substances which exist in this kind of a structure are Na₂O, K₂O and Rb₂O.

Cesium Halide Structure

 Chloride ions are primitive cubic while the cesium ion occupies the center of the unit cell. There is one chloride ion and one cesium ion per unit cell. Therefore the formula is CsCl. The coordination number of cesium is eight and that of chloride ions is also eight. Other substances which exist in this kind of a structure are all halides of cesium.

Corundum Structure

The general formula of compounds crystallizing in corundum structure is A₂O₃^. The closest packing is that of anions (oxide) in hexagonal primitive lattice and two-third of the octahedral voids are filled with trivalent cations. Examples are : Fe₂O₃, Al₂O₃ and Cr₂O₃.

Pervoskite Structure

The general formula is ABO₃. One of the cation is bivalent and the other is tetravalent. Example: CaTiO₃, BaTiO₃. The bivalent ions are present in primitive cubic lattice with oxide ions on the centers of all the six square faces. The tetravalent cation is in the center of the unit cell occupying octahedral void.

Spinel and Inverse Spinel Structure

Spinel is a mineral (MgAl₂O₄). Generally they can be represented as M²⁺M₂³⁺O₄ , where M²⁺ is present in one-eighth of tetrahedral voids in a FCC lattice of oxide ions and M³⁺ ions are present in half of the octahedral voids. M²⁺ is usually Mg, Fe, Co, Ni, Zn and Mn; M³⁺ is generally Al, Fe, Mn, Cr and Rh. Examples are ZnAl₂O₄^, Fe₃O₄^, FeCr₂O₄ etc. Many substances of the type also have this structure. In an inverse spinel the ccp is of oxide ions, M²⁺ is in one-eight of the tetrahedral voids while M³⁺ would be in one-eight of the terahedral voids and one-fourth of the octahedral voids.

Exercise 3: The unit cell of silver iodide (AgI) has 4 iodine atoms in it. How many silver atoms must be there in the unit cell.

Exercise 4: The co-ordination number of the barium ions, Ba²⁺, in barium chloride (BaF₂) is 8. What must be the co-ordination number of the fluoride ions, .

Imperfections in a Crystal

The discovery of imperfections in an other wise ideally perfect crystal is one of the most fascinating aspects of solid state science. An ideally perfect crystal is one which has the same unit cell and contains the same lattice points throughout the crystal. The term imperfection or defect is generally used to describe any deviation of the ideally perfect crystal from the periodic arrangement of its constituents.

Point Defect

If the deviation occurs because of missing atoms, displaced atoms or extra atoms, the imperfection is named as a point defect. Such defects can be the result of imperfect packing during the original crystallisation or they may arise from thermal vibrations of atoms at elevated temperatures because with increase in thermal energy there is increased probability of individual atoms jumping out of their positions of lowest energy. The most common point defects are the Schottky defect and the Frenkel defect. Comparatively less common point defects are the metal excess defect and the metal deficiency defects. All these defects have been discussed below in some details.

Schottky Defects: The defects arise if some of the lattice points in a crystal are unoccupied. The points which are unoccupied are called lattice vacancies. The existence of two vacancies, one due to a missing Na⁺ ion and the other due to a missing Cl ion in a crystal of NaCl, is shown in figure 16. The crystal, as a whole remains neutral because the number of missing positive and negative ions is the same.

Figure 16

Schottky defects appear generally in ionic crystals in which the positive and the negative ions do not differ much in size. Sodium chloride and cesium chloride furnish good examples of ionic crystals in which Schottky defects occurs.

Frenkel Defects:

These defects arise when an ion occupies an interstitial position between the lattice points. This is shown in figure for the crystal of AgBr.

As can be seen, one of the Ag⁺ ions occupies a position in the interstitial space rather than its own appropriate site in the lattice. A vacancy is thus created in the lattice as shown. It may be noted again that the crystal remains neutral since the number of positive ions is the same as the number of negative ions. The presence of Ag⁺ ions in the interstitial space of AgBr crystal is responsible for the formation of a photographic image on exposure of AgBr crystals (i.e., photographic plate) to light.

ZnS is another crystal in which Frenkel defects appear. Zn²⁺ ions are entrapped in the interstitial space leaving vacancies in the lattice.

Frenkel defects appear in crystals in which the negative ions are much larger than the positive ions. Like Schottky defects, the Frenkel defects are also responsible for the conduction of electricity in crystals and also for the phenomenon of diffusion in solids.

Metal Excess Defects. The Colour Centres. It has been observed that if a crystal of NaCl is heated in sodium vapour, it acquires a yellow colour. This yellow colour is due to the formation of a non-stoichiometric compound of sodium chloride in which there is a slight excess of sodium ions. What happens in this case is that some sodium metal gets doped into sodium chloride crystal which, due to the crystal energy, gets ionised into Na⁺ and e. This electron occupies a site that would otherwise be filled by a chloride ion, as illustrated in figure.

Figure 18

There is evidently an excess of metal ions although the crystal as a whole is neutral. A little reflection would show that there are six Na⁺ sites adjacent to the vacant site occupied by the electron. The extra electron in thus shared between all the six Na⁺ ions which implies that this electron is not localised at the vacant Cl site. On the other hand, this electron is similar to the delocalised p electrons present in molecules containing conjugate double bonds. Light is absorbed when this delocalised electron makes an easy transition from its ground state to an excited state. As a result, the non stoichiometric form of sodium chloride appears coloured. Because of this, the sites occupied by the extra electrons are known as colour centres. These are also called F-centres. This name comes from the German word Farbe meaning colour. The non-schiometric sodium chloride may be represented by the formula Na_(1+d)Cl where d is the excess sodium metal doped in the crystal because of its exposure to sodium vapour.

Another common example of metal excess defects is the formation of a magenta coloured non-stoichiometric compound of potassium chloride by exposing the crystals of KCl to K metal vapour. The coloured compound contains an excess of K⁺ ions, the vacant Cl sites being filled by electrons obtained by the ionization of the excess K metal doped in to the crystal.

Metal Deficiency Defects. In certains cases, one of the positive ions is missing from its lattice site and the extra negative charge is balanced by some nearby metal ion acquiring two charges instead of one. There is evidently, a deficiency of the metal ions although the crystal as a whole is neutral. This type of defect is generally found amongst the compounds of transition metals which can exhibit variable valency. Crystals of FeO, FeS and NiO show this type of defects. The existence of metal deficiency defects in the crystal of FeO is illustrated.

It is evident from the above discussion that all types of point defects result in the creation of vacancies or `holes' in the lattice of the crystals. The presence of holes lowers the density as well as the lattice energy or the stability of the crystals. The presence of too many holes may cause a partial collapse of the lattice.

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Magnetic Properties of Solids - The Solid State, Class 12, Chemistry

Magnetic properties :

The macroscopic magnetic properties of material are consequence of magnetic moments associated with the individual electron. Each electron in an atom has magnetic moment due to two reasons. The first one is due to the orbital motion around the nucleus, and the second is due to the spin of electron around its own axis. A moving electron may be considered to be a small current loop, generating a small magnetic field, and having a magnetic moment along its axis of rotation.

The other type of magnetic moment originates from electron spin which is directed along the spin axis. Each electron in an atom may be thought of as being a small magnet having permanent orbital and spin magnetic moments. The fundamental magnetic moment is Bohr Magneton, m_s equal to 9.27 × 10^-24 Am² . For each electron in an atom, the spin magnetic moment is ± m_B (depending upon the two possiblities of the spin). The contribution of the orbital magnetic moment is equal to m_i m_B, where m_i is the magnetic quantum number of the electron.

Based on the bahaviour in the external magnetic field, the substances are divided different catagories.

(i) Diamagnetic substances : Substances which are weakly repelled by the external magnetic field are called diamagnetic substances, e.g. TiO₂, NaCl, Benzene.

(ii) Paramagnetic substances : Substances which are attracted by the external magnetic field are called paramagnetic substances, e.g., O₂, Cu²⁺, Fe³⁺.

(iii) Ferromagnetic substances : Substances which show permanent magnetsim even in the absence of the magnetic field are called ferromagnetic substances e.g., Fe, Ni, Co, CrO₂.

(iv) Ferrimagnetic Substances : Substances which are expected to possess large magnetism on the basis of the unpaired electrons but actually have small net magnetic moment are called ferrimagnetic substances, e.g., Fe₃O₄.

(v) Anti-ferromagnetic substances : Substances which are expected to posses paramagnetic behaviour, ferromagnetism on the basis of unpaired electron but actually they possess zero not magnetic moment are called anti-ferromagnetic substances, e.g., MnO₂

Note : Ferromagnetic, Anti-ferromagnetic and Ferrimagnetic solids change into paramagnetic at some temperature. It may be further pointed out here that each ferromagnetic substance has a characteristic temperature above which no ferromagnetism is observed. This is known as Curie Temperature.

Dielectric Properties :

Insulators do not conduct electricity because the electrons present in them are hel tightly to the individual atoms or ions and are not free to move. however, when electric field is applied, polarization takes place because nuclei are attracted to one side and the electron cloud to the other side. Thus dipoles are formed. In addition to these dipoles, there may also be permanent dipoles in the crystal. These dipoles may align themselves in an ordered manner so that such crystals have a net dipole moment.

Such polar crystals show the following interersting propeties :

(i) Piezoelectricity or Pressure - electricity : When mechanical stress is applied on such crystals so as to deform them, electricity is produced due to displacement of ions. The electricity thus produced is called piezoelectricity and the crystals are called piezoelectric crystals. A few examples of piezoelectric crystals include Titanates of Barium and Lead, Lead zirconate (PbZrO₃), Ammonium dihydrogen phosphate (NH₄H₂PO₄) and Quarts. These crystals are used as pick-ups in record players where they prodduce electrical signals by application of pressure. They are also used in microphones, ultrasonic generators and sonar detectors.

(ii) Ferroelectricity : In some of the piezoelectric crystals, the dipoles are permanently polarized even in the absence of the electric field. However, on applying, electric field, the direction of polarization changes, this phenomenon is called ferroelectricity due to analogy with ferromagnetism. Some examples of the ferroelectric solids are barium titanate (BaTiO₃), Sodium potassium tartarate (Rochelle salt) and Potassium dihydrogen phosphate (KH₂PO₄)

Note : All ferroelectric solids are piezoelectric but the reverse is not true.

(iii) Anti-ferroelectricity : In some crystals, the dipoles align themselves in such a way that alternately, they point up and down so that the crystal does not posses any net dipole moment. Such crystals are said to be anti-ferroelectric. A typical example of such crystals is Lead zirconate (PbZrO₃)

Superconductivity :

A substance is said to be superconducting when it offers no resistance to the flow of electricity.

Isomorphism :

The property shown by crystals of different chemical substances which exhibit the same crystalline form is known as isomorphism.

Such substances have similar chemical formula, e.g., Na₂HPO₄. 12H₂O and Na₃ASO₄. 12H₂O

Chemically substances having same type of formulae are not necessarily Isomorphous.

Magnetic Properties of Crystals

Electrical Properties of Crystals