Class 12 Chemistry Ch1 The Solid State Notes PDF Download

The chapter, Solid State sounds like a full-on crystal maze at first, but honestly, once you crack the basics, it’s actually super chill and totally scoring. This chapter is deleted from the latest CBSE Class 12 Syllabus.

These Class 12 Chemistry The Solid State Notes have everything you need: NCERT facts, crystal-clear definitions, key formulas, and short tricks to remember packing types without the panic. Neat, no-nonsense, and exam-friendly.

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Class 12 Chemistry Ch 1 Notes 

Whether you need clear definitions, solved numericals, shortcut tips, or just a solid understanding of what's actually going on in the chapter - it's all here. Scroll once, and you’re sorted.

What are Solids and their Types?

A solid is a state of matter where the particles are tightly packed, locked in place, and can only vibrate. Because of this strong force of attraction, solids have a definite shape, fixed volume, high density, and are usually rigid. You can’t squish them, and they don’t flow like liquids or gases.

Types of Solids (Based on Particle Arrangement)

Solids are classified into two major types depending on how their particles are arranged:

1. Crystalline Solids

Think of these as solids with discipline. Their particles are arranged in a perfect, repeating pattern that extends throughout the structure - like a neatly laid brick wall.

Features:

• Long-range order in particle arrangement
• Sharp, fixed melting point
• Different properties in different directions (anisotropy)
• Cut into clean, flat surfaces
• Considered true solids

Examples: Sodium chloride (NaCl), diamond, quartz, ice

Also Check out our CBSE Class 12 PYQs with detailed solutions.

2. Amorphous Solids

Now these are the “go-with-the-flow” kind of solids. Particles are arranged randomly, without any fixed long-range pattern - more like blobs stacked on each other.

Features:

• No regular arrangement (short-range order only)
• Melt over a temperature range (not sharp)
• Same properties in all directions (isotropy)
• Break into uneven, irregular pieces
• Often called pseudo-solids or supercooled liquids

Examples: Glass, plastic, rubber

Classification of Crystalline Solids

So now that we know what crystalline solids are, here’s the deal - not all of them are the same inside. Their internal structure depends on what kind of particles they’re made of and what kind of forces hold those particles together.

Here are the four main types of crystalline solids you need to remember:

1. Ionic Solids

These are made of positive and negative ions, like Na⁺ and Cl⁻ in table salt. They’re held together by strong electrostatic forces, which is why they’re usually hard and have high melting points.

But here's a catch - they don’t conduct electricity in solid form. Once you melt them or dissolve them in water, they do.

Examples: Sodium chloride (NaCl), magnesium oxide (MgO), potassium chloride (KCl)

2. Molecular Solids

These are formed by neutral molecules - either polar, non-polar, or hydrogen-bonded. Since they’re held by weak forces (like van der Waals or hydrogen bonding), they’re generally soft, have low melting points, and are non-conductors.

They’re further divided into:

• Non-polar (like H₂, Cl₂)
• Polar (like SO₂)
• Hydrogen-bonded (like ice, NH₃)

Examples: Ice, dry ice (solid CO₂), iodine

3. Covalent or Network Solids

These are like one giant molecule where every atom is bonded to another through strong covalent bonds. That makes them extremely hard, with very high melting points, and most of them don’t conduct electricity.

Examples: Diamond, silicon carbide (SiC), quartz (SiO₂)

Exception: Graphite conducts electricity, even though it’s covalent - thanks to free electrons between its layers.

4. Metallic Solids

These are made entirely of metal atoms, with a unique bonding style. The metal atoms release electrons, which float around like a “sea” of electrons. That’s why metals are good conductors, shiny, malleable, and can be drawn into wires.

Examples: Copper, iron, silver, aluminium

Crystal Lattice and Unit Cell

Let's see how solids are arranged inside. When we talk about crystalline solids, we’re basically talking about a regular pattern in which particles are stacked. This repeating pattern is called a crystal lattice.

What is a Crystal Lattice?

A crystal lattice is a 3D arrangement of points showing how particles (atoms, ions, or molecules) are arranged in a solid. It’s like a never-ending grid - the same structure repeats again and again throughout the solid.

There are seven basic types of crystal systems (like cubic, tetragonal, hexagonal, etc.) — but you don’t need to go super deep into them unless specified.

What is a Unit Cell?

Imagine cutting a piece of the lattice that shows the entire pattern - just smaller. That tiny portion is called a unit cell. A unit cell is the smallest repeating unit of a crystal lattice that, when repeated in space, creates the whole structure of the solid.

So basically:

Unit cell = one brick
Crystal lattice = the whole wall made from that brick

Types of Unit Cells

There are two main types:

1. Primitive unit cell – particles are only present at the corners
2. Centered unit cells – apart from corners, particles are also present at:
   
   
   • Body-centered → one in the center
   • Face-centered → one at each face
   • End-centered → one on two opposite faces only

Number of Atoms per Unit Cell

This part is all about how many atoms are actually present in one unit cell - depending on its type.

In crystalline solids, atoms are arranged in a specific repeating structure called the unit cell. However, atoms are not always completely inside one unit cell - many are shared with neighboring unit cells. 

So, when we calculate the number of atoms per unit cell, we only consider the fraction of each atom that actually belongs to that particular unit cell.

This depends on the position of the atom in the unit cell.

Contribution of Atoms Based on Position:

• Corner atom → shared by 8 unit cells → contributes 1/8
• Face-centered atom → shared by 2 unit cells → contributes 1/2
• Body-centered atom → fully inside the cell → contributes 1

(Edge-centered atoms contribute 1/4, but are not common in standard cubic cells)

Number of Atoms in Different Cubic Unit Cells:

1. Simple Cubic Unit Cell (Primitive):

• Atoms only at the 8 corners
• 8 corners × 1/8 = 1 atom per unit cell

2. Body-centered Cubic (BCC):

• Atoms at 8 corners + 1 atom at the body center
• (8 × 1/8) + 1 = 2 atoms per unit cell

3. Face-centered Cubic (FCC):

• Atoms at 8 corners + 6 face centers
• (8 × 1/8) + (6 × 1/2) = 1 + 3 = 4 atoms per unit cell

Packing Efficiency and Voids in Solids

In solid-state chemistry, not all the space in a crystal is used by the particles (atoms, ions, or molecules). Some part of the space remains empty (voids). The part that is actually filled by particles is called packing efficiency.

So basically:

Packing Efficiency = How much of the unit cell is filled
Voids = The empty space left in the unit cell

What is Packing Efficiency?

Packing efficiency is the percentage of the total space in a unit cell that is occupied by the constituent particles.

Packing Efficiency=(Volume occupied by particlesTotal volume of unit cell)×100\text{Packing Efficiency} = \left( \frac{\text{Volume occupied by particles}}{\text{Total volume of unit cell}} \right) \times 100Packing Efficiency=(Total volume of unit cellVolume occupied by particles​)×100

Packing Efficiency in Different Unit Cells

1. Simple Cubic (SC):

• Only corner atoms
• Total atoms per unit cell = 1
• Packing efficiency = 52.4%Voids (empty space) = 47.6%
• Least efficient — not very dense

2. Body-Centered Cubic (BCC):

• Corner atoms + one in the center
• Total atoms per unit cell = 2
• Packing efficiency = 68%
• Voids = 32%
• Found in iron and other common metals

3. Face-Centered Cubic (FCC) or Cubic Close Packing (CCP):

• Corner atoms + face-centered atoms
• Total atoms per unit cell = 4
  Packing efficiency = 74% (highest among all)
• Voids = 26%
• Most efficient arrangement
• Seen in metals like copper, silver, aluminium

What are Voids?

Voids are the unoccupied spaces inside the unit cell. They’re important in determining the porosity, stability, and density of a solid.

More packing → less voids → higher densities

Students should also try solving a few related questions from Class 12 Chemistry Sample Papers for better retention.

Types of Defects in Solids

Solids are not always perfect. When some atoms or ions are missing, added, or misplaced in the crystal, it’s called a defect. These defects can affect density, colour, and conductivity.

Let’s look at the important types:

1. Stoichiometric Defects

The formula (ratio of ions) stays correct.

There are 2 kinds:

• Vacancy defect:
  Some atoms or ions are missing from their place.
  → The solid becomes lighter (lower density)
  
  
• Interstitial defect:
  An extra atom or ion is stuck in a gap (not a regular position).
  → The solid becomes heavier (higher density)

Example: These defects can happen in NaCl or simple solids.

2. Impurity Defect

Something foreign is added to the crystal.

• Sometimes we add another element into a solid.
• This creates vacancies to keep charges balanced.
• For example: Adding SrCl₂ to NaCl
  → Sr²⁺ replaces 2 Na⁺ → creates a gap (cation vacancy)

Helps in doping semiconductors.

3. Non-Stoichiometric Defects

The formula of the compound changes (wrong ratio).

There are 2 types:

• Metal Excess Defect:
  A negative ion is missing, but an extra electron is left behind.
  → The solid conducts electricity & may look coloured
  → Example: NaCl turns yellow, ZnO turns yellow when heated
• Metal Deficiency Defect:
  A positive ion is missing, and nearby ones change charge to balance
  → Example: FeO, where Fe²⁺ is missing and replaced by Fe³⁺

Electrical and Magnetic Properties of Solids

Solids don’t just sit there - some of them conduct electricity, some don’t. Some even show magnetism. This depends on how the electrons move and how ions or atoms are arranged. Let’s break it down:

Electrical Properties of Solids

Solids are divided into 3 types based on electrical conductivity:

1. Conductors
   
   • Metals like Cu, Ag
   • Have free electrons that move easily
   • Conduct heat and electricity well
2. Insulators
   
   • Example: Glass, rubber
   • No free electrons → do not conduct electricity
3. Semiconductors
   
   • Example: Silicon (Si), Germanium (Ge)
   • Conduct electricity only under special conditions
   • Used in electronics and chips

Doping in Semiconductors

To improve conductivity, we add impurities. This is called doping.

• n-type semiconductor:
  Extra electrons are added
  → e.g., Si + Phosphorus (P)
  
  
• p-type semiconductor:
  Creates holes (missing electrons)
  → e.g., Si + Boron (B)
  
  

Magnetic Properties of Solids

Based on how atoms respond to magnetic fields, there are 5 types:

1. Diamagnetic
   
   • Weakly repelled by a magnetic field
   • Example: Zn, H₂O
2. Paramagnetic
   
   • Weakly attracted
   • Have unpaired electrons
   • Example: O₂, Cu²⁺
3. Ferromagnetic
   
   • Strongly attracted
   • Have unpaired electrons aligned in one direction
   • Example: Fe, Co, Ni
4. Antiferromagnetic
   
   • Equal but opposite magnetic moments cancel out
   • Net magnetism = 0
   • Example: MnO
5. Ferrimagnetic
   
   • Unequal magnetic moments
   • Partial cancellation → some magnetism
   • Example: Fe₃O₄, MgFe₂O₄

Conclusion

And that’s a wrap on The Solid State - from tricky unit cells to simple crystal logic, you’ve made it through. This chapter builds the base for so many others, so getting it clear now is a smart move.

If these The Solid State Class 12 Notes PDF helped even a bit, give yourself a small win - it counts. One chapter done, and you’re already ahead of the rush.

FAQs

Q1. What are the four main types of crystalline solids?

Ans. They are molecular, ionic, metallic, and network (covalent) solids.

Q2. What are Schottky and Frenkel defects?

Ans.

• Schottky defect: equal numbers of cations and anions are missing, keeping the formula unchanged.
• Frenkel defect: a cation leaves its place and occupies a nearby gap, without changing the formula.

Q3. Why are amorphous solids isotropic?

Ans. Amorphous solids have no long-range order, so their physical properties are the same in all directions.

Q4. Why does ZnO turn yellow on heating?

Ans. Heating ZnO creates metal-excess defects - electrons trapped in anion vacancies absorb visible light, causing a yellow colour.

Q5. How does doping make semiconductors more conductive?

Ans. Doping adds impurities that create extra electrons (n-type) or holes (p-type), boosting electrical conductivity.