Class 12 Chemistry Notes Chapter 1 (The solid state) – Chemistry-I Book

Alright class, let's get straight into Chapter 1: The Solid State. This is a fundamental chapter, and understanding the structure and properties of solids is crucial for many concepts in chemistry and material science, making it important for your government exam preparations. Pay close attention to the definitions, classifications, calculations, and defects.

Chapter 1: The Solid State - Detailed Notes for Exam Preparation

1. Introduction to Solids:

• Definition: Solids are states of matter characterized by definite shape, volume, and high density. They are rigid and incompressible.
• Reason for Rigidity: Constituent particles (atoms, molecules, or ions) are held in fixed positions by strong intermolecular forces and can only oscillate about their mean positions.

2. Classification of Solids:
Based on the arrangement of constituent particles:

• (a) Crystalline Solids:
  • Arrangement: Particles are arranged in a definite, ordered, repeating three-dimensional pattern (long-range order).
  • Properties:
    • Sharp melting point.
    • Anisotropic (physical properties like electrical resistance or refractive index show different values when measured along different directions in the same crystal).
    • Have a definite enthalpy of fusion.
    • Give clean cleavage when cut.
    • Examples: NaCl, Quartz, Diamond, Sugar, Metals (Fe, Cu, Ag).
• (b) Amorphous Solids:
  • Arrangement: Particles are arranged randomly, lacking a regular pattern (short-range order only, if any). Structure is similar to supercooled liquids.
  • Properties:
    • Melt over a range of temperatures.
    • Isotropic (physical properties are the same in all directions).
    • Do not have a definite enthalpy of fusion.
    • Give irregular cleavage when cut.
    • Sometimes called pseudo-solids or supercooled liquids.
    • Examples: Glass, Rubber, Plastics, Tar, Amorphous Silicon.

3. Classification of Crystalline Solids (Based on Bonding):

4. Crystal Lattices and Unit Cells:

• Crystal Lattice (Space Lattice): A regular three-dimensional arrangement of points in space representing the constituent particles (atoms, ions, or molecules).
• Lattice Points: The positions occupied by the constituent particles in the crystal lattice.
• Unit Cell: The smallest repeating portion of a crystal lattice which, when repeated in different directions, generates the entire lattice.
• Parameters of a Unit Cell: Characterized by:
  • Edge lengths: a, b, c
  • Angles between edges: α (between b & c), β (between a & c), γ (between a & b).
• Types of Unit Cells:
  • Primitive (Simple): Particles only at the corners.
  • Centred: Particles at corners plus other positions:
    • Body-Centred (BCC): Particles at corners + one particle at the body center.
    • Face-Centred (FCC): Particles at corners + one particle at the center of each face.
    • End-Centred: Particles at corners + one particle at the center of any two opposite faces.

5. Seven Crystal Systems: Based on the unit cell parameters (a, b, c and α, β, γ). (Memorize the names and conditions, especially for Cubic).

6. Calculation of Number of Atoms in a Unit Cell (Z):

• Contribution of particle at:
  • Corner: 1/8
  • Face Centre: 1/2
  • Body Centre: 1
  • Edge Centre: 1/4
• Simple Cubic (SC): 8 corners × (1/8) = 1 atom (Z=1)
• Body-Centred Cubic (BCC): (8 corners × 1/8) + (1 body centre × 1) = 1 + 1 = 2 atoms (Z=2)
• Face-Centred Cubic (FCC): (8 corners × 1/8) + (6 face centres × 1/2) = 1 + 3 = 4 atoms (Z=4)

7. Close Packing in Solids:
Arrangement of spheres (particles) to minimize empty space.

• 1D: Spheres in a row (Coordination Number = 2)

• 2D:

  • Square Close Packing (AAA type): Spheres in adjacent rows placed directly below each other. Coordination Number = 4. Packing Efficiency = 52.4%.
  • Hexagonal Close Packing (ABAB type): Spheres in adjacent rows placed in the depressions of the first row. Coordination Number = 6. Packing Efficiency = 60.4%. More efficient than square close packing.

• 3D: Stacking of 2D layers.

  • From 2D Square Close Packed layers: Forms Simple Cubic lattice (AAA stacking).
  • From 2D Hexagonal Close Packed layers:
    • ABAB... Stacking (Hexagonal Close Packing - hcp): Second layer spheres over depressions of the first, third layer spheres vertically above the first layer. Coordination Number = 12. Packing Efficiency = 74%. Examples: Mg, Zn.
    • ABCABC... Stacking (Cubic Close Packing - ccp or Face-Centred Cubic - FCC): Second layer spheres over depressions of the first, third layer spheres over the depressions not covered by the second layer. Fourth layer aligns with the first. Coordination Number = 12. Packing Efficiency = 74%. Examples: Cu, Ag, Au.

• Coordination Number (CN): The number of nearest neighbours of a particle in a crystal lattice.

  • SC: 6
  • BCC: 8
  • hcp/ccp (FCC): 12

8. Packing Efficiency:
The percentage of total space filled by the particles.

• Packing Efficiency = (Volume occupied by spheres in the unit cell / Total volume of the unit cell) × 100
• SC: 52.4%
• BCC: 68%
• hcp/ccp (FCC): 74% (Most efficient)

9. Voids (Interstitial Sites):
Empty spaces left between the spheres in close-packed structures.

• Tetrahedral Void: Surrounded by 4 spheres. Smaller void.
• Octahedral Void: Surrounded by 6 spheres. Larger void.
• Relationship: If the number of close-packed spheres = N, then:
  • Number of Octahedral Voids = N
  • Number of Tetrahedral Voids = 2N

10. Calculations Involving Unit Cell Dimensions:
Density (ρ) of a crystal:
ρ = (Z × M) / (a³ × N<0xE2><0x82><0x90>)
Where:

• Z = Number of atoms per unit cell
• M = Molar mass (g/mol)
• a = Edge length of the unit cell (usually in cm or pm; 1 pm = 10⁻¹⁰ cm)
• N<0xE2><0x82><0x90> = Avogadro's number (6.022 × 10²³ mol⁻¹)
• a³ = Volume of the unit cell (for cubic)

11. Imperfections in Solids (Crystal Defects):
Deviations from the perfectly ordered arrangement.

• (a) Point Defects: Irregularities around a point or an atom.

  • I. Stoichiometric Defects: Do not disturb the stoichiometry. Also called intrinsic or thermodynamic defects.
    • Vacancy Defect: A lattice site is vacant. Decreases density. Common in non-ionic solids. Arises due to heating.
    • Interstitial Defect: Some constituent particles occupy interstitial sites. Increases density. Common in non-ionic solids.
    • Frenkel Defect (Dislocation Defect): An ion (usually cation, being smaller) leaves its lattice site and occupies an interstitial site. Density remains unchanged. Shown by ionic solids with a large difference in ion sizes. Examples: AgCl, AgBr, AgI, ZnS.
    • Schottky Defect: Equal numbers of cations and anions are missing from their lattice sites to maintain electrical neutrality. Decreases density. Shown by ionic solids with similar cation/anion sizes and high coordination numbers. Examples: NaCl, KCl, CsCl, AgBr. (Note: AgBr shows both Frenkel and Schottky defects).
  • II. Non-Stoichiometric Defects: Disturb the stoichiometry.
    • Metal Excess Defect:
      • Due to Anionic Vacancies: An anion is missing, and the electron occupies the site (F-centre) to maintain neutrality. F-centres (Farbenzenter) impart colour (e.g., NaCl heated in Na vapour becomes yellow). Leads to paramagnetism.
      • Due to Extra Cations at Interstitial Sites: An extra cation occupies an interstitial site, and an electron occupies another interstitial site for neutrality. Example: ZnO heated becomes yellow (Zn²⁺ + ½O₂ + 2e⁻).
    • Metal Deficiency Defect: A cation is missing, and the charge is balanced by an adjacent metal ion having a higher positive charge. Common in transition metal compounds showing variable oxidation states. Example: FeO (often found as Fe₀.₉₅O).
  • III. Impurity Defects: Foreign atoms are present at lattice sites or interstitial sites.
    • Example: Doping SrCl₂ into NaCl. Sr²⁺ replaces two Na⁺ ions, occupying one site and leaving one cation vacancy. Increases conductivity.

• (b) Line Defects: Irregularities along rows of lattice points (e.g., dislocations). (Less focus at 12th level).

12. Electrical Properties:
Based on conductivity:

• Conductors: High conductivity (10⁴ to 10⁷ ohm⁻¹m⁻¹). Metals. Conductivity decreases with increasing temperature.
• Insulators: Very low conductivity (10⁻²⁰ to 10⁻¹⁰ ohm⁻¹m⁻¹). Non-metals, molecular solids.
• Semiconductors: Intermediate conductivity (10⁻⁶ to 10⁴ ohm⁻¹m⁻¹). Conductivity increases with increasing temperature. Examples: Si, Ge.
  • Intrinsic Semiconductors: Pure semiconductors (Si, Ge). Conductivity is low but increases with temperature.
  • Extrinsic Semiconductors (Doping): Conductivity increased by adding impurities.
    • n-type: Doping Group 14 (Si, Ge) with Group 15 element (P, As). Excess electrons are charge carriers.
    • p-type: Doping Group 14 (Si, Ge) with Group 13 element (B, Al). Electron holes (vacancies) are charge carriers.

13. Magnetic Properties:
Based on behaviour in a magnetic field:

• Diamagnetic: Weakly repelled by the magnetic field. All electrons are paired. Examples: H₂O, NaCl, C₆H₆.
• Paramagnetic: Weakly attracted by the magnetic field. Have unpaired electrons. Lose magnetism in the absence of the field. Examples: O₂, Cu²⁺, Fe³⁺, Cr³⁺.
• Ferromagnetic: Strongly attracted by the magnetic field. Can be permanently magnetized. Domains align parallel in the field. Examples: Fe, Co, Ni, Gd, CrO₂.
• Antiferromagnetic: Domains align anti-parallel, cancelling out the magnetic moment. Example: MnO.
• Ferrimagnetic: Domains align parallel and anti-parallel in unequal numbers, resulting in a net magnetic moment. Weakly attracted compared to ferromagnetic. Example: Fe₃O₄ (Magnetite), Ferrites (MgFe₂O₄, ZnFe₂O₄).

Note: Ferromagnetic, Antiferromagnetic, and Ferrimagnetic substances become Paramagnetic above a certain temperature called the Curie Temperature (T<0xE1><0xB5><0x9C>).

Practice MCQs for Exam Preparation:

1. Which of the following is an amorphous solid?
   (a) Graphite
   (b) Quartz glass (SiO₂)
   (c) Chrome alum
   (d) Silicon carbide (SiC)

2. The number of atoms per unit cell in a Body-Centred Cubic (BCC) structure is:
   (a) 1
   (b) 2
   (c) 4
   (d) 6

3. The packing efficiency of a Face-Centred Cubic (FCC) structure is:
   (a) 52.4%
   (b) 68%
   (c) 74%
   (d) 90%

4. Schottky defect in crystals is observed when:
   (a) Density of the crystal is increased.
   (b) Unequal number of cations and anions are missing from the lattice.
   (c) An ion leaves its normal site and occupies an interstitial site.
   (d) Equal number of cations and anions are missing from the lattice.

5. Doping Silicon (Group 14) with Phosphorus (Group 15) results in:
   (a) p-type semiconductor
   (b) n-type semiconductor
   (c) Metal conductor
   (d) Insulator

6. In a close-packed structure of N spheres, the number of tetrahedral voids is:
   (a) N/2
   (b) N
   (c) 2N
   (d) 4N

7. Which type of crystalline solid is characterized by low melting point, softness, and poor electrical conductivity?
   (a) Ionic
   (b) Metallic
   (c) Covalent/Network
   (d) Molecular

8. An element crystallizes in a BCC lattice with a cell edge length of 288 pm. The radius of the atom is: (For BCC, √3a = 4r)
   (a) (√3/4) × 288 pm
   (b) (√2/4) × 288 pm
   (c) (1/2) × 288 pm
   (d) (1/4) × 288 pm

9. Which of the following exhibits antiferromagnetism?
   (a) Fe₃O₄
   (b) MnO
   (c) Ni
   (d) TiO₂

10. F-centres in ionic crystals are:
    (a) Cation vacancies
    (b) Anion vacancies occupied by unpaired electrons
    (c) Interstitial cations
    (d) Frenkel defects

Answer Key for MCQs:

1. (b) Quartz glass (SiO₂) - Note: Quartz is crystalline, Quartz glass is amorphous.
2. (b) 2
3. (c) 74%
4. (d) Equal number of cations and anions are missing from the lattice.
5. (b) n-type semiconductor
6. (c) 2N
7. (d) Molecular
8. (a) (√3/4) × 288 pm
9. (b) MnO
10. (b) Anion vacancies occupied by unpaired electrons

Make sure you revise these concepts thoroughly. Focus on the definitions, differences between types of solids and defects, coordination numbers, packing efficiencies, and the density calculation formula. Good luck with your preparation!