Solid State Chemistry Class 12 Handwritten Notes PDF Download | Class 12 chemistry ch-1 solid state notes

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Solid State Chemistry Class 12 Handwritten Notes PDF Download | Class 12 chemistry ch-1 solid state notes
Solid State Chemistry Class 12 Handwritten Notes PDF Download | Class 12 chemistry ch-1 solid state notes

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Solid State Chemistry Class 12 Handwritten Notes PDF Download :-

Class 12 chemistry notes of solid state in details :-

Introduction

Apart from liquid and gaseous states, solid state is a state of matter. Solids have very strong intermolecular interactions, and there are very few vacant spaces between the atoms/ions/molecules. As a result, they have a predetermined shape and volume.

Characteristic Properties of Solids

The following properties come under the category of solids:

  • Solids have high density.
  • Solids have low compressibility.
  • Solids are rigid in nature.
  • Solids have definite shape and volume.

Classification of Solids

On the basis of the following parameter, solids are broadly classified as:

  • Classification based on various properties.
  • Classification based on bonding present in building blocks.

On the Basis of Various Properties

On the basis of the various properties of solids, they can be classified as:

  • Crystalline solids
  • Amorphous solids

Amorphous solids have an uneven structure over long distances and lack sharp properties, whereas crystalline solids have a regular structure throughout the entire volume and sharp qualities. The table below shows the many differences.

PropertyCrystalline solidsAmorphous Solids
ShapeCrystalline solids have a long range order. Amorphous solids have a short range order.
Melting pointCrystalline solids tend to have definite melting points.Amorphous solids do not have a definite melting point.
Heat of fusionCrystalline solids have a definite heat of fusion.Amorphous solids do not have a definite heat of fusion.
CompressibilityCrystalline solids are rigid and incompressible.Amorphous solids may be compressed to some extent.
Cutting with a sharp edged toolCrystalline solids tend to break into two pieces with plane surfaces.Amorphous solids give irregular cleavage, that means they break into two pieces with irregular shape.
Isotropy and AnisotropyCrystalline solids are anisotropic.Amorphous solids are isotropic.
Volume changeWhen crystalline solids melt, there is a sudden change in their volume.On melting there is no sudden change in the volume of amorphous solids.
SymmetryCrystalline solids possess symmetry.Amorphous solids do not possess any symmetry.
Interfacial anglesCrystalline solids have interfacial angles.Amorphous solids do not have interfacial angles.

Based on Bonding

Solids are classified according to the sort of bonding present in their building units. The table below lists many types of solids as well as their properties. 

The Different Properties of the Four Types of Solids are Listed as:

Type of SolidConstituentBondingExamplesPhysicalElectricalMelting
Molecular SolidsNon Polar
MoleculesDispersion or London forcesArAr, CCl4CCl4 ,H2H2 ,I2I2 , CO2CO2 SoftInsulatorVery low
PolarMoleculesDipole-Dipole interactionsHCl, SO2HCl, SO2 SoftInsulatorLow
Hydrogen BondedMoleculesHydrogen bondingH2OH2O (ice)HardInsulatorLow
Ionic SolidsIonsCoulombic or electrostaticNaCl, MgO,NaCl, MgO, ZnS, CaF2 ZnS, CaF2 Hard but brittleInsulators in solid state but conductors in molten state and in aqueous solutions.High
Metallic SolidsPositive ions in a sea of delocalised electrons.Metallic bondingFe, Cu,Fe, Cu,Ag, MgAg, MgHard but malleable and ductileConductors in solid state as well as in molten state.Fairly high
Covalent or network solidsAtomsCovalent bondingSiO2SiO2 (quartz)SiC, CSiC, C (diamond)C(graphite)C(graphite) Hard

Soft
Insulators

Conductor
Very High

Structure of Crystalline Solids

Crystal Lattice and Unit Cell

The crystalline solid regular array of building pieces (atoms/ions/molecules) is known as the “Crystal Lattice.”

“Unit Cell” refers to the smallest component of a crystal lattice that can be repeated in all directions to form the full crystal lattice.

Small spheres represent the atoms of ions or molecules in a unit cell. Variations in the following parameters produce several lattices:

  • The edge along the 3 axis – a, b, c.
  • The interfacial angle – α,β,γα,β,γ 
  • Location of atoms/ions with respect to each other in crystal lattice.

(Image will be uploaded soon)

Primitive Unit Cells and Bravais Lattices

There are seven different types of unit cells, as well as various subtypes of unit cells. Primitive Unit Cells or Crystal Habits are the names given to these seven unit cells. The following are listed in the table below:

Crystal SystemAxial DistanceAxial anglesExamples
Cubica=b=ca=b=c α=β=γ=90∘α=β=γ=90∘ Copper, Zinc blende, KClKCl 
Tetragonala=b≠ca=b≠c α=β=γ=90∘α=β=γ=90∘White tin, SnO2,TiO2SnO2,TiO2 
Orthorhombica≠b≠ca≠b≠c α=β=γ=90∘α=β=γ=90∘Rhombic sulphur, CaCO3CaCO3 
Monoclinica≠b≠ca≠b≠cα=γ=90∘α=γ=90∘ ; β≠90∘β≠90∘ Monoclinic sulphur, PbCrO2PbCrO2 
Hexagonala=b≠ca=b≠c α=β=90∘α=β=90∘ ; γ=120∘γ=120∘ Graphite, ZnOZnO 
Rhombohedrala=b=ca=b=c α=β=γ≠90∘α=β=γ≠90∘ Calcite (CaCO3)(CaCO3) Cinnabar (HgS)(HgS) 
Triclinica≠b≠ca≠b≠c α≠β≠γ≠90∘α≠β≠γ≠90∘ K2Cr2O7, CuSO4.5H2OK2Cr2O7, CuSO4.5H2O 

For these 7 types of unit cells, 14 types of Lattices exist in nature. These 14 lattices are named as “Bravais Lattices”.

Crystal SystemSpace LatticeExamples
Cubic a=b=ca=b=cHere a, b and c are the dimensions of a unit cell along three axes. α=β=γ=90∘α=β=γ=90∘ Here, αα, ββ and γγ are the sizes of three angles between the axes. Simple: Lattice points at the eight corners of the unit cells.(Image will be uploaded soon)Body Centered: Points at the eight corners and at the body centre.(Image will be uploaded soon)Face Centered: Points at the eight corners and at the six face enters.(Image will be uploaded soon)Pb, Hg, AgPb, Hg, AgAu, Cu, ZnSAu, Cu, ZnS Diamond, KClKClNaClNaCl, Cu2OCu2O , CaF2CaF2 and alumns, etc.  
Tetragonala=b≠ca=b≠c α=β=γ=90∘α=β=γ=90∘Simple: Points at the eight corners of the unit cell.(Image will be uploaded soon)Body Centered: Points at the eight corners and at the body centre. (Image will be uploaded soon)SnO2,TiO2,SnO2,TiO2,  ZnO2,NiSO4ZnO2,NiSO4  ZrSiO4,ZrSiO4,  PbWO4PbWO4And white tin.
Orthorhombic:a≠b≠ca≠b≠c α=β=γ=90∘α=β=γ=90∘ Simple: Points at the eight corners of the unit cell.(Image will be uploaded soon)End Centered: Also called side centered or base centered. Points at the eight corners and at two face centers opposite to each other.(Image will be uploaded soon)Body Centered: Points at the eight corners and at the body centre.(Image will be uploaded soon)Face Centered: Points at the eight corners and at the six face centres.(Image will be uploaded soon)KNO3, K2SO4,KNO3, K2SO4,  PbCO3, BaSO4PbCO3, BaSO4Rhombic sulphur, MgSO4.7H2OMgSO4.7H2O etc.
Rhombohedral or Trigonal a=b=ca=b=c ,α=β=γ≠90∘α=β=γ≠90∘Simple: Points at the edge corners of the unit cell.(Image will be uploaded soon)NaNO3, CaSO4NaNO3, CaSO4 , calcite, quartz, As, Sb, BiAs, Sb, Bi 
Hexagonala=b≠ca=b≠c ,α=β=90∘α=β=90∘  γ=120∘γ=120∘Simple: Points at the twelve or points at the twelve corners of the unit cell out corners of the hexagonal lined by thick line, prism and at the centres of top and bottom faces.(Image will be uploaded soon)ZnO, PbS, CdSZnO, PbS, CdS , graphite, ice, Mg, Zn, CdMg, Zn, Cd etc.
Monoclinica≠b≠ca≠b≠c α=γ=90∘α=γ=90∘ , β≠90∘β≠90∘ Simple: Points at the eight corners of the unit cell.(Image will be uploaded soon)End Centered: Point at the eight corners and two face centres opposite to each other.(Image will be uploaded soon)Na2SO4.10H2ONa2SO4.10H2O , Na2B4O7.10H2ONa2B4O7.10H2O , CaSO4.2H2OCaSO4.2H2O , monoclinic sulphur etc.
Triclinica≠b≠ca≠b≠c α≠β≠γ≠90∘α≠β≠γ≠90∘ Simple: Points at eight corners of the unit cell.(Image will be uploaded soon)CaSO4.5H2OCaSO4.5H2O , K2Cr2O7K2Cr2O7 , H3BO3H3BO3 

The focus will primarily be on cubic unit cells and their arrangements.

Cubic Unit Cells

The most common unit cell is this one. The atoms or spheres in a cubic unit cell can be found at the following locations.

  • Corners
  • Body centre
  • Face centres

The contributions of a sphere stored at various locations are as follows:

LocationContribution
Corners1/8
Body Centre1
Face Centre1/2

Types of Cubic Unit Cells

(Image will be uploaded soon)

The following factors distinguish these unit cells from one another:

  • The positions of the spheres within the unit cell.
  • The unit cell’s rank (effective number of spheres inside a unit cell).
  • The relationship between the radius and the edge length of a single sphere.
  • Fractional packing (fraction of volume occupied by spheres in a unit cell).

The following parameters are provided in the table below for all three unit cells:

Type of Cubic CrystalNo. of atoms at different locationsStructureRankPackingRelation b/w atomic radius and edge length (a)
CornersBody CentresFace Centre
Simple Cubic8(Image will be uploaded soon)152%r = a/2
Body Centred81(Image will be uploaded soon)268%r=3a−−√4r=3a4 
Face Centred86(Image will be uploaded soon)474%r=2a−−√4r=2a4 

Density of Cubic Crystals

By the following formula, the density of the cubic crystal is given:

\rho =M×Za3×NA\rho =M×Za3×NA 

Where, Z is the rank of the unit cell, M is the molar mass of the solid, a is the edge length of the unit cell, NA is the Avogadro number.

The volume of Z will depend on the type of unit cell.

Close Packing in Solids: Origin of Unit Cells

Assume we have a set of spheres of identical size that we must arrange in a single layer with the requirement that the spheres be in close proximity to one another. There are two sorts of layers that can be used:

  • Square Packing
  • Hexagonal Packing

Spheres are arranged in square packing in such a way that the rows are both horizontal and vertical. The Co-ordination number is 4 in this situation.

(Image will be uploaded soon)

It is more efficient to pack hexagonally. It has a Coordination number of 6 and has fewer voids than square packing. 

If we add another layer to the square packing, we can do the following:

  • A comparable layer is placed just above the foundation layer, with the second layer’s spheres appearing just above the first layer’s spheres, and the layers are repeated. If the first layer is designated as A, the packing is of the type AA, and the unit cell is simple cubic.
  • On the other hand, we get BCC unit cells and ABAB type of packing when spheres from the second layer are inserted in depressions from the first layer.

The following are examples of hexagonal foundation layer arrangements:

When we place the second hexagonal layer A in the depressions of the first hexagonal layer A, we get two sorts of voids. Hollow and through voids of layer A and layer B are the X kind of voids. Layer B voids that are directly above spheres in layer A are referred to as Y type voids. When the spheres of the second layer are placed over Y voids, layer 1 is repeated, and ABABAB type packing is obtained. The hexagonal unit cell is obtained in this arrangement, and the packing is known as hexagonal close packing (HCP). This packing has a 74 percent efficiency.

When the third layer is applied to X-type voids, a new layer C is created, and the process is repeated. Packing of the ABCABCABC type will be obtained. The FCC unit cell is used in this design, and the packing efficiency is 74%.

(Image will be uploaded soon)

(Image will be uploaded soon)

Voids

Definition

Voids are the empty spaces inside a sphere. The amount and shape of voids is determined by the unit cell and packing used.

Radius Ratio

The radius ratio of a sphere that can be perfectly fit in the void to the radius of surrounding spheres is used to determine the size of the void. This is written as:

Radius ratio = rRrR 

Types of Voids

Trigonal Void

It is the void formed of equal radii which touches each other as shown in the figure.

FigureKey Points
(Image will be uploaded soon)Radius RatiorR=0.155rR=0.155 Smallest voidCoordination number is 3.

Tetrahedral Void

It is formed by the contact of four spheres and is located in the centre of a tetrahedron formed by the contact of four spheres.

FigureKey Points
(Image will be uploaded soon)Radius ratio rR=0.225rR=0.225 Number of voids in FCC crystals is 8.Position at a distance: a3–√4a34 from every corner.Coordination number is 4. 

Octahedral Void

FigureKey Points
(Image will be uploaded soon)Radius ratio rR=0.414rR=0.414 Number of voids in FCC crystals is 4.Positions: Body centre and edge centre.Rank is 4.Coordination number is 6.

Cubic Void

The voids formed by the close contact of eight spheres.

The following are the key points:

  • Radius ratio is equal to rR=0.732rR=0.732 
  • Number of voids in a cubic crystal is 1.
  • Position is at the body centre.
  • Coordination number is 8.
  • Rank 1.

It is clear from the above details that:

Trigonal < Tetrahedral < Octahedral < Cubic

Classification of Ionic Structures

The simultaneous arrangement of cations and anions in a lattice/unit cell produces ionic compounds. The larger of two species takes up major places in a unit cell, while the lesser species takes up vacancies in proportion to their size. Which is determined by the radius ratio. Below is a list of the various ratios.

Limiting Radius Ratiox=r + r – x=r + r – C. N. ShapeExample 
x<0.155x<0.155 2LinearBeF3BeF3 
0.155⩽x⩽0.2250.155⩽x⩽0.225 3Planar TriangularAlCl3AlCl3 
0.225⩽x⩽0.4140.225⩽x⩽0.414 4TetrahedronZnSZnS 
0.414⩽x⩽0.7320.414⩽x⩽0.732 6OctahedronNaClNaCl 
0.732⩽x⩽0.9990.732⩽x⩽0.999 8Body centred cubicCsClCsCl 

On the basis of these ratio ranges, the ionic crystal is classified into five categories which are as follows:

NaClNaCl Type Structure

FigureKey Points
(Image will be uploaded soon)Rock salt structureCl−Cl− occupy corners and face centres and Na + Na +  occupy octahedral voids in FCC crystal.Effective formula is Na4Cl4Na4Cl4 Coordination number of Na + Na + is 6.Coordination number of Cl−Cl−is 6.Distance b/w the nearest neighbour [rNa + +rCl−=a2][rNa + +rCl−=a2] 

ZnSZnS Type Structure

FigureKey Points
(Image will be uploaded soon)Zinc Blende StructureS2−S2− ions occupy main positions and Zn + 2Zn + 2 ions are present in alternate tetrahedral voids in FCC crystal.Effective formula is Zn4S4Zn4S4 The Coordination of Zn + 2Zn + 2is 4.Coordination number of S2−S2−is 4.rZn + 2+rS2−=a3–√4rZn + 2+rS2−=a34 

Fluorite Type Structure

FigureKey Points
(Image will be uploaded soon)Ca + 2Ca + 2 ions occupy main positions and F−F− ions occupy tetrahedral voids in FCC crystal.Effective formula is Ca4F8Ca4F8 .The Coordination number of Ca + 2Ca + 2 is 8.Coordination number of F−F−is 4.rZn + 2+rS2−=a3–√4rZn + 2+rS2−=a34

Anti Fluorite Structure

FigureKey Points
(Image will be uploaded soon)Anti-fluorite structureCommon in alkali oxides like Na2O, Li2ONa2O, Li2O etc.O2−O2− ions occupy FCC and Li + Li +  ions occupy the tetrahedral voids.The Coordination number of Li + Li + is 4.Coordination number of O2−O2−is 8.rLi++rO2−=a3–√4rLi++rO2−=a34

CsClCsCl Type Structure

FigureKey Points
(Image will be uploaded soon)Cesium Halide structureCl−Cl− ions simple cubic locations (corners) and Cs + Cs +  ions occupy body centre in BCC lattice.Effective lattice isCsClCsCl.The Coordination number of Cs + Cs + is 8.Coordination number of Cl−Cl−is 8.rCs++rCl−=a3–√4rCs++rCl−=a34

Imperfections in Solids

In a crystal structure sometimes some imperfections or defects occur:

Classification of Defects

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Vacancies

This type of defect occurs when the positions that should contain atoms or ions are vacant.

Interstitial Sites

These are sites located between regular positions; sometimes atoms or ions may occupy these positions.

(Image will be uploaded soon)

Stoichiometric Defects

The stoichiometry of solids are not disturbed by these defects.

Schottky Defects

In ionic solids, it’s a vacancy defect. Electrical neutrality is maintained because the number of missing cations and anions is equal. The density of the substance is reduced as a result of this flaw. Ionic compounds with almost identical cation and anion sizes demonstrate the flaw. Examples are: KCl, NaCl, AgBrKCl, NaCl, AgBr etc.

(Image will be uploaded soon)

Frenkel Defect

The smaller ion is relocated from its typical position to an interstitial region in ionic solids. At its original place, it causes a vacancy defect, and at new locations, it causes an interstitial defect. Dislocation defect is another name for it. It has no effect on the solid’s density. Ionic compounds with a considerable disparity in ion size are examples of this type of defect.

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Note: Silver bromide (AgBr)(AgBr) shows Schottky and Frenkel defects both.

Non Stoichiometric Defects

The compounds with these flaws have combining components in a different ratio than their stoichiometric formulas require.

Metal Excess Defect

Due to Anionic Vacancies: It’s possible that the anion is absent from its lattice position, leaving an electron behind to keep the charge balanced. The F centre is the electron-containing site. They provide the crystal colour; F stands for Farbenzenter, which means colour. This defect looks like schottky defect and can be seen in crystals with schottky defect. Examples: NaCl, KClNaCl, KCl etc.

(Image will be uploaded soon)

Due to the Presence of Extra Cations in the Interstitial Sites.

To maintain electrical neutrality, an additional cation may be present in one interstitial site while an electron is present in another interstitial site. This is a flaw that is similar to the Frenkel defect and can be discovered in crystals with the Frenkel defect.

(Image will be uploaded soon)

Metal Deficieny Defect

When metal has a fluctuating valency, this is a defect. FeOFeO , for example, is generally found in compositions ranging from Fe0.93OFe0.93O  to Fe0.96OFe0.96O . Some Fe + 2Fe + 2  cations are missing from FeOFeO  crystals, but the loss of positive charge is compensated for by the existence of the requisite amount of Fe + 3Fe + 3  ions.

Revision Notes of Chemistry Class 12 Chapter 1- The Solid State

Notes of Chemistry Class 12 Chapter 1

Class 12 Chapter 1 Chemistry The Solid State introduces students to the basic concept of Solids, their types and characteristics. To understand all the essential concepts covered in Chemistry Chapter 1, students first understand what is meant by solids. The definition of Solids, along with its types and other crucial concepts related to it are covered in the well-explained and straightforward way in Class 12th Chemistry Chapter 1 notes.

Solids: Such chemical substances which have definite size, shape, volume and rigidity are known as Solid substances. These substances have high density and low compressibility.

There Are Two Types of Solids

  • Amorphous Solids
  • Crystalline Solids

Other Crucial Concepts Covered Under Class 12 Chemistry Ch 1 Notes:

  • Classification of Solids
  • Structure of Crystalline Solids
  • Different types of Solids
  • Square Packing 
  • Hexagonal Packing 
  • Types of Cubic Unit Cells 
  • Voids
  • Types of Voids
  • Trigonal Voids
  • Octahedral Voids
  • Cubic Voids
  • Ionic Structure 
  • Fluorite Structure 
  • Metal Excess Defect
  • Metal Deficiency Effect

Some Important Questions of Chapter 1 Chemistry Class 12 

1. Define the term ‘amorphous’. Give a few examples of amorphous solids.

2. What makes a glass different from a solid such as quartz? Under what conditions could quartz be converted into glass?

3. How many e lattice points are there in one unit sale of each of the following lattices? 

  • Face centred cubic
  • Face centred tetragonal
  • Body-centred

4. Define the structure of crystalline solids.

5. Define the types of cubic unit cells.

6. Classify the ionic structures.

7. Explain the metal deficiency defect

Benefits of Using Class 12 Chemistry Chapter 1 Revision Notes

To make learning fun and easy for students, Class 12 Chemistry Chapter 1 notes are available in PDF file to download for free from the official website of Vedantu and its app. After downloading, students can access these notes offline as well for the quick revision of essential terms and concepts.

CBSE Class 12 Chemistry Chapter 1 notes will help the students in practising the questions that are going to be asked in their examinations. These notes will help them score better in their board examinations.

Conclusion

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