Thermal properties of matter 2

Thermal Properties of Matter - Class 11 Notes

1. Heat and Temperature

1.1 Basic Concepts

Heat

Form of energy transferred between systems due to temperature difference

SI unit: Joule (J)

Other common unit: Calorie (1 cal = 4.186 J)

Temperature

Measure of hotness or coldness of a body

Determines direction of heat flow (from higher to lower temperature)

SI unit: Kelvin (K)

Other scales: Celsius (°C), Fahrenheit (°F)

Temperature Conversion Formulas

K = °C + 273.15

°F = (9/5)°C + 32

°C = (5/9)(°F - 32)

NEET Tip

Remember: Temperature is a scalar quantity while heat transfer has direction (vector-like but not a vector). Temperature determines thermal equilibrium.

1.2 Thermal Equilibrium

Definition

When two bodies are in thermal contact and there is no net heat flow between them

At thermal equilibrium, both bodies have the same temperature

Zeroth Law of Thermodynamics

If two systems A and B are each in thermal equilibrium with a third system C, then A and B are in thermal equilibrium with each other

This law establishes temperature as a fundamental property

2. Thermal Expansion

2.1 Linear Expansion

Formula

ΔL = L₀αΔT

Where: ΔL = change in length, L₀ = original length, α = coefficient of linear expansion, ΔT = change in temperature

Final length: L = L₀(1 + αΔT)

Coefficient of Linear Expansion (α)

Defined as fractional change in length per degree temperature change

Unit: K⁻¹ or °C⁻¹

For most solids: α ≈ 10⁻⁵ to 10⁻⁶ K⁻¹

2.2 Area Expansion

Formula

ΔA = A₀βΔT

Where: β = coefficient of area expansion ≈ 2α

Final area: A = A₀(1 + βΔT)

2.3 Volume Expansion

Formula

ΔV = V₀γΔT

Where: γ = coefficient of volume expansion ≈ 3α

For liquids: γ is much larger than for solids

Final volume: V = V₀(1 + γΔT)

Anomalous Expansion of Water

Water contracts when heated from 0°C to 4°C (density increases)

Above 4°C, water expands normally

Maximum density of water is at 4°C (1 g/cm³)

Important for aquatic life in cold climates

Bihar Board Focus

Numerical problems on thermal expansion are frequently asked. Practice problems involving bimetallic strips, gaps in railway tracks, and pendulum clocks.

3. Calorimetry

3.1 Specific Heat Capacity

Definition

Amount of heat required to raise the temperature of unit mass of a substance by 1°C (or 1K)

SI unit: J kg⁻¹ K⁻¹

Formula

Q = mcΔT

Where: Q = heat supplied, m = mass, c = specific heat capacity, ΔT = temperature change

Specific Heat Capacities of Common Substances

Substance	Specific Heat Capacity (J g⁻¹ K⁻¹)
Water	4.186
Ice	2.09
Aluminium	0.897
Iron	0.449
Copper	0.385

3.2 Heat Capacity

Definition

Amount of heat required to raise the temperature of a body by 1°C

C = mc (where m is mass, c is specific heat capacity)

SI unit: J K⁻¹

3.3 Principle of Calorimetry

Principle

When two bodies at different temperatures are brought into thermal contact, heat flows from hotter to colder body until thermal equilibrium is reached

Heat lost by hotter body = Heat gained by colder body (assuming no heat loss to surroundings)

Calorimetry Equation

m₁c₁(T₁ - T) = m₂c₂(T - T₂) + ...

Where T is the final equilibrium temperature

NEET Tip

Calorimetry problems often involve mixing of substances at different temperatures. Remember to account for phase changes if they occur during the process.

4. Change of State

4.1 Latent Heat

Definition

Amount of heat required to change the state of unit mass of a substance at constant temperature

Two types: Latent heat of fusion (solid to liquid) and Latent heat of vaporization (liquid to gas)

Formula

Q = mL

Where: Q = heat supplied, m = mass, L = latent heat

Unit of L: J kg⁻¹

Latent Heat Values

Substance	Latent Heat of Fusion (kJ/kg)	Latent Heat of Vaporization (kJ/kg)
Water	334	2260
Lead	23	870
Mercury	11.3	295

4.2 Phase Diagrams

Triple Point

Temperature and pressure at which all three phases (solid, liquid, gas) coexist in equilibrium