Physics · Class 9 · CBSE Chapter 11

Heat and Temperature

Understanding thermal energy, its measurement, transfer, and effects on matter

−273.15°C Absolute zero (0 K)

100°C Boiling point of water

4186 J Specific heat of water (per kg·°C)

3 Modes of heat transfer

Heat vs Temperature — The Core Distinction

Heat is a form of energy that flows from a hotter body to a cooler body due to a temperature difference. Temperature is a measure of the average kinetic energy of the particles in a substance — it tells us how hot or cold a body is, not how much thermal energy it contains.

Analogy

Think of a swimming pool and a cup of tea. The tea is at a higher temperature, but the pool contains far more heat energy because it has vastly more water molecules. Temperature is the 'intensity'; heat is the 'quantity'.

A large iceberg has more total heat energy than a small cup of boiling water — even though the cup is at a much higher temperature. This is why heat and temperature are fundamentally different quantities.

Heat (Q) is measured in Joules (J) or Calories (cal) — it is energy
Temperature (T) is measured in Celsius (°C), Fahrenheit (°F), or Kelvin (K)
Heat flows spontaneously only from higher temperature to lower temperature
Two bodies in thermal equilibrium have the same temperature — no net heat flows between them
The SI unit of heat is the Joule; 1 calorie = 4.186 J

Temperature Scale Conversions

\[ K = °C + 273.15 \qquad °F = \frac{9}{5}\,°C + 32 \]

K

Temperature in Kelvin — the SI base unit. Always positive; 0 K is absolute zero.

°C

Celsius scale — water freezes at 0°C, boils at 100°C at 1 atm.

°F

Fahrenheit scale — water freezes at 32°F, boils at 212°F.

Note

Kelvin and Celsius have the same size degree; only the zero point differs. Always convert to Kelvin for gas law calculations.

K

Kelvin

SI base unit of temperature

°C

Celsius

Common scientific scale

°F

Fahrenheit

Common in USA

Temperature Scales Comparison

Celsius, Fahrenheit and Kelvin scales side by side — showing key reference points

Heat Energy Formula

\[ Q = m c \Delta T \]

Q

Heat energy absorbed or released (Joules, J)

m

Mass of the substance (kilograms, kg)

c

Specific heat capacity — the heat needed to raise 1 kg by 1°C (J/kg·°C)

ΔT

Change in temperature = Final temperature − Initial temperature (°C or K)

Note

If Q is positive, the body absorbs heat (heats up). If Q is negative, the body releases heat (cools down).

Q

Heat energy

Joules (J)

m

Mass

kilograms (kg)

c

Specific heat capacity

J / (kg·°C)

ΔT

Temperature change

°C or K (same magnitude)

Heat Calculator — Q = mcΔT

Adjust mass, specific heat, and temperature change to calculate heat energy. Try comparing water (c = 4186) with iron (c = 450) for the same mass and temperature change.

\( Q = m \times c \times \Delta T \)

m

1 kg

Mass (kg)

c

4186 J/kg°C

Specific heat (J/kg·°C)

ΔT

20 °C

Temperature change (°C)

Heat energy Q

— J

Specific Heat Capacities of Common Substances

4186

Water

J/kg·°C — highest of common substances; why oceans moderate climate

2090

Ice

J/kg·°C — half that of liquid water

900

Aluminium

J/kg·°C — heats and cools quickly

490

Iron / Steel

J/kg·°C — used in cookware

385

Copper

J/kg·°C — good thermal conductor

128

Mercury

J/kg·°C — lowest among listed liquids

Three Modes of Heat Transfer

1

Conduction Solids mainly

Heat transfer through direct contact — molecule to molecule. Faster-moving (hotter) molecules collide with slower (cooler) neighbours, transferring kinetic energy. Occurs primarily in solids. Metals are excellent conductors because free electrons carry energy rapidly.

2

Convection Fluids (liquid + gas)

Heat transfer by actual movement of a fluid (liquid or gas). Hot fluid becomes less dense, rises; cooler fluid sinks to replace it — forming a convection current. This is how room heaters warm a room, how the atmosphere circulates, and how oceanic currents form.

3

Radiation No medium needed

Heat transfer by electromagnetic waves (infrared radiation) — requires no medium whatsoever. All objects above absolute zero emit radiation. The Sun heats Earth entirely by radiation across the vacuum of space. Dark, rough surfaces absorb and emit radiation better than light, shiny surfaces.

Three Modes of Heat Transfer

Conduction (left) through a metal rod, Convection (centre) in a liquid, Radiation (right) from a hot body to a cold body through vacuum

Conduction vs Convection vs Radiation

Property	Conduction	Convection	Radiation
Medium required	Yes — solid preferred	Yes — fluid (liquid/gas)	No — works in vacuum
How energy moves	Molecular vibration	Bulk fluid movement	Electromagnetic waves (IR)
Speed	Slow in non-metals, fast in metals	Moderate	Speed of light (3×10⁸ m/s)
Best materials	Metals (copper, aluminium)	Water, air	Dark, rough surfaces absorb best
Real example	Iron rod heated at one end	Room heater warming a room	Sun warming Earth
Prevention	Thermal insulators (wood, foam)	Vacuum flask, double glazing	Reflective/shiny surfaces

Latent Heat — Heat Without Temperature Change

Latent heat is the heat energy absorbed or released by a substance during a change of state (solid↔liquid or liquid↔gas) at constant temperature. The temperature does not change during the phase transition even though heat is being added or removed.

Analogy

Imagine melting ice — you keep adding heat but the temperature stays at 0°C until all the ice has melted. The energy goes into breaking the bonds between water molecules, not into raising the temperature.

This is why steam burns are more dangerous than boiling water burns at the same 100°C — steam releases an additional 2.26 MJ/kg of latent heat when it condenses on your skin.

Latent heat of fusion (L_f) — heat for solid↔liquid change at melting point
Latent heat of vaporisation (L_v) — heat for liquid↔gas change at boiling point
For water: L_f = 334,000 J/kg and L_v = 2,260,000 J/kg
During a phase change, all added heat goes into changing the state, not the temperature
Formula: Q = mL, where L is the specific latent heat

Latent Heat Formula

\[ Q = m L \]

Q

Heat energy absorbed or released during phase change (Joules)

m

Mass of the substance undergoing the phase change (kg)

L

Specific latent heat — heat per unit mass for the phase change (J/kg)

L_f

Specific latent heat of FUSION (melting/freezing): water = 3.34 × 10⁵ J/kg

L_v

Specific latent heat of VAPORISATION (boiling/condensing): water = 2.26 × 10⁶ J/kg

L_f (water)

Latent heat of fusion

3.34 × 10⁵ J/kg

L_v (water)

Latent heat of vaporisation

2.26 × 10⁶ J/kg

Heating Curve of Water

Temperature vs heat added for water — flat regions show phase changes where temperature stays constant despite heat being added

Key Formulas Summary

Formula	What it calculates	When to use
Q = mcΔT	Heat energy for temperature change	When state does NOT change — just heating or cooling
Q = mL	Heat energy for phase change	When state CHANGES — melting, boiling, freezing, condensing
K = °C + 273	Convert Celsius to Kelvin	Gas law problems; absolute temperature needed
°F = (9/5)°C + 32	Convert Celsius to Fahrenheit	When Fahrenheit value is asked
ΔT = T_f − T_i	Calculate temperature change	Finding Q using Q = mcΔT

Test Your Understanding

Q1 100 g of water is heated from 20°C to 70°C. How much heat is absorbed? (c_water = 4200 J/kg°C)

A. 21 J

B. 2100 J

C. 21000 J

D. 210 J

Q = mcΔT = 0.1 × 4200 × (70−20) = 0.1 × 4200 × 50 = 21000 J. Remember: mass must be in kg, so 100 g = 0.1 kg.

Q2 Which mode of heat transfer does NOT require any medium?

A. Conduction

B. Convection

C. Radiation

D. All require a medium

Radiation transfers heat as electromagnetic (infrared) waves and can travel through vacuum — this is how the Sun's heat reaches Earth. Conduction and convection both require a medium.

Q3 Convert 37°C (normal body temperature) to Kelvin.

A. 37 K

B. 236 K

C. 310 K

D. 373 K

K = °C + 273 = 37 + 273 = 310 K. This is an important value to remember — normal human body temperature is 310 K.

Q4 A block of ice at 0°C is being heated. What happens to its temperature as it melts?

A. It rises steadily from 0°C

B. It stays at 0°C until all ice melts

C. It first drops then rises

D. It rises to 100°C immediately

During a phase change, ALL the heat energy goes into breaking molecular bonds, not increasing kinetic energy. So temperature remains constant at 0°C until the entire ice has melted. This is latent heat.

Q5 Why does a steel spoon feel colder than a wooden spoon at the same room temperature?

A. Steel has a lower temperature than wood

B. Steel conducts heat away from your hand faster than wood

C. Wood stores more heat than steel

D. Steel has higher specific heat capacity

Both objects are at the same room temperature. Steel is a much better thermal conductor than wood — it draws heat away from your hand rapidly, making it feel cold. Your hand loses heat faster to the steel spoon.

Q6 How much heat is needed to melt 500 g of ice at 0°C? (L_f = 3.36 × 10⁵ J/kg)

A. 168,000 J

B. 16,800 J

C. 1680 J

D. 336,000 J

Q = mL = 0.5 × 3.36 × 10⁵ = 168,000 J = 168 kJ. Note: 500 g = 0.5 kg. This energy goes purely into melting, not raising temperature.

💡

Did you know?

Water has one of the highest specific heat capacities of any common substance (4186 J/kg°C). This is why coastal cities have milder climates than inland cities — the ocean absorbs enormous amounts of heat in summer and releases it slowly in winter, moderating temperature extremes.

✗

Common Mistake

Students often confuse heat and temperature. A large cold lake has more total thermal energy (heat) than a small hot cup of tea. Temperature measures intensity (average KE per molecule); heat measures total energy. They are NOT the same quantity.

✓

Exam Tip — Which formula to use?

Q = mcΔT when temperature CHANGES (no phase change). Q = mL when temperature stays CONSTANT (phase change occurs). In a multi-step problem, you may need BOTH: first heat ice to 0°C (Q=mcΔT), then melt it (Q=mL), then heat water to 100°C (Q=mcΔT again).

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Units Alert

Always convert grams to kilograms before using Q = mcΔT or Q = mL. The values of c and L are given per kilogram. Using grams gives an answer 1000 times too small — the most common calculation error in this chapter.

Experiment: Specific Heat Capacity of a Metal

Aim: To determine the specific heat capacity of a metal using the method of mixtures

Materials

Metal block (aluminium or iron) of known mass
Boiling water and heat source
Calorimeter (insulated container) with known specific heat
Known mass of cold water
Thermometer (0–110°C)
Electronic balance
Tongs

Procedure

Record the mass of the metal block (m_metal), calorimeter (m_cal), and cold water (m_w).
Heat the metal block in boiling water until its temperature equals that of boiling water (T_hot ≈ 100°C). Record this as T_hot.
Measure and record the initial temperature of cold water + calorimeter (T_cold).
Using tongs, quickly transfer the hot metal block into the cold water in the calorimeter. Stir gently.
Record the maximum temperature reached by the mixture (T_mix).
Apply the principle of conservation of energy: Heat lost by metal = Heat gained by water + Heat gained by calorimeter.

Observation

The temperature of the metal drops from T_hot to T_mix. The temperature of the water rises from T_cold to T_mix. The final equilibrium temperature T_mix lies between T_cold and T_hot.

Conclusion

Using: m_metal × c_metal × (T_hot − T_mix) = m_w × c_w × (T_mix − T_cold) + m_cal × c_cal × (T_mix − T_cold), we can solve for c_metal. The experiment demonstrates conservation of thermal energy and the concept of thermal equilibrium.

History of Heat and Temperature Science

1593

Galileo Galilei

Invented the thermoscope — the first device to show temperature changes, though without a numerical scale

1714

Gabriel Fahrenheit

Developed mercury thermometer and established the Fahrenheit scale with reproducible fixed points

1742

Anders Celsius

Proposed the Celsius (centigrade) scale with 0° at boiling and 100° at freezing — inverted by Linnaeus to today's convention

1798

Count Rumford

Showed that mechanical work can generate unlimited heat (boring cannon barrels), disproving the caloric theory

1824

Sadi Carnot

Founded thermodynamics; showed that heat engines cannot be 100% efficient

1848

Lord Kelvin

Defined absolute zero and established the Kelvin temperature scale based on thermodynamic principles

1850

Rudolf Clausius

Formulated the Second Law of Thermodynamics — heat flows spontaneously only from hot to cold

1865

Rudolf Clausius

Introduced the concept of entropy — a measure of disorder in a system

Key Vocabulary

Heat

Old English: haetu (warmth, heat)

Energy in transit due to a temperature difference. Flows from higher to lower temperature. Measured in Joules.

Q = mcΔT = 2 × 4200 × 10 = 84,000 J of heat absorbed

Temperature

Latin: temperare (to mix, moderate)

Measure of average kinetic energy of molecules in a substance. Determines direction of heat flow.

37°C = 310 K = 98.6°F — normal human body temperature

Specific heat capacity

Latin: species (kind) + calere (to be warm)

Heat energy required to raise the temperature of 1 kg of a substance by 1°C. Symbol: c, unit: J/kg°C.

c_water = 4186 J/kg°C; c_iron = 490 J/kg°C

Latent heat

Latin: latere (to lie hidden) — heat is 'hidden' as it doesn't change temperature

Heat absorbed or released during a phase change at constant temperature. From Latin: latens (hidden).

L_f(ice) = 3.34×10⁵ J/kg; L_v(water) = 2.26×10⁶ J/kg

Conduction

Latin: conducere (to lead together)

Heat transfer through a material by molecular vibration and collision, without bulk movement of the material.

A metal spoon in hot soup — handle gets warm by conduction

Convection

Latin: convehere (to carry together)

Heat transfer by bulk movement of a heated fluid (liquid or gas) due to density differences.

Land and sea breezes; hot air balloon rising

Radiation

Latin: radiare (to emit beams)

Heat transfer by electromagnetic waves (infrared) — requires no medium and travels at the speed of light.

Warmth felt from a campfire across distance, solar heating

Thermal equilibrium

Greek: thermos (hot) + Latin: aequilibrium (equal balance)

State when two bodies in contact reach the same temperature and net heat flow between them stops.

Coffee left in a room eventually reaches room temperature

Fascinating Heat & Temperature Facts

01

The specific latent heat of vaporisation of water (2.26 MJ/kg) is nearly 7 times its latent heat of fusion (0.334 MJ/kg) — which is why boiling requires far more energy than melting.

02

At absolute zero (0 K or −273.15°C), molecular motion theoretically ceases completely. It has never been achieved experimentally — only approached to within billionths of a degree.

03

The human body generates roughly 80 watts of heat at rest — about the same as a standard incandescent light bulb. During intense exercise this rises to 1000+ watts.

04

In a vacuum flask (Thermos), all three modes of heat transfer are minimised: the vacuum prevents conduction/convection, and the silvered glass reflects radiated heat.

05

Lightning superheats the surrounding air to approximately 30,000 K — about 5 times hotter than the surface of the Sun (5,778 K).

06

The oceans store about 1000 times more heat than the atmosphere — which is why they are the primary driver of Earth's long-term climate rather than air temperature alone.

Putting on a sweater doesn't make you warm — it just slows how fast the heat your body generates escapes. The sweater has no heat of its own.

— Thermodynamics perspective

Technically, there is no such thing as 'cold' — only the absence of heat. Cold is not a substance; it is just low thermal energy.

— Physics of heat

If you blew on a hot cup of tea to cool it, you used convection. If you held it near a window on a cold day, you used radiation. Physics is just cooking with extra steps.

— Applied thermodynamics

Applications of Heat Transfer in Daily Life

1

Vacuum Flask (Thermos)

Minimises all three heat transfers: vacuum eliminates conduction/convection; silvered surfaces reduce radiation. Keeps liquids hot or cold for hours.
2

Cooking vessels

Copper and aluminium pots used for fast, even heat distribution (high conductivity). Handles made of wood or plastic to prevent burning hands (low conductivity).
3

Central heating systems

Hot water heated by a boiler circulates through radiators by convection — heats rooms by radiation from the radiator surface and convection of room air.
4

Greenhouse effect

Earth's atmosphere lets in short-wave solar radiation but traps long-wave infrared re-radiated from the surface — same principle as a glass greenhouse.
5

Refrigerator

Removes heat from inside by evaporating a refrigerant (latent heat of vaporisation absorbs heat) then compressing it elsewhere to release that heat outside.
6

Steam engine / Power plant

Uses latent heat of vaporisation to store large amounts of energy in steam. Converting water to steam absorbs 2.26 MJ/kg — released as mechanical work on turbines.

Common Doubts

Why does sweating cool the body?

Sweat (water) evaporates from the skin surface. Evaporation requires latent heat of vaporisation (about 2.4 MJ/kg). This latent heat is drawn from the skin and body, cooling you down. The same principle explains why wet clothes feel cold — the water evaporates and takes heat from your body.

Why is the temperature of a mixture always between the temperatures of the two components?

Heat flows from the hotter body to the cooler body until thermal equilibrium is reached. The hotter object loses energy (cools), the cooler one gains energy (warms). Conservation of energy means the final temperature must lie between the two starting temperatures — it cannot exceed the hotter or drop below the cooler initial temperature.

Why does ice at 0°C feel colder than water at 0°C?

Both are at the same temperature (0°C). However, when ice touches your skin, it first needs to absorb latent heat of fusion (334,000 J/kg) before it can melt — this heat comes from your hand, cooling it further. Water at 0°C only needs to absorb sensible heat (Q = mcΔT) to warm up, which is much less per unit mass.

What is the difference between heat capacity and specific heat capacity?

Heat capacity (C) is the heat needed to raise an entire object's temperature by 1°C — it depends on both the material AND the mass. Specific heat capacity (c) is the heat needed per kilogram — it depends only on the material. Relationship: C = mc. A big iron block has a large heat capacity; iron itself has a specific heat of 490 J/kg°C regardless of size.

Why is the Kelvin scale important in science?

Kelvin starts at absolute zero — the point of minimum possible thermal energy. Unlike Celsius (which can go negative), Kelvin is always positive and directly proportional to average molecular kinetic energy. This makes it essential for gas laws (PV = nRT uses T in Kelvin), thermodynamic equations, and any calculation where a ratio of temperatures is needed.

Chapter Summary — Heat and Temperature

Class 9 Physics · CBSE Chapter 11 · Complete revision guide

Key Takeaways

Heat is energy (Joules); Temperature is a measure of average molecular KE (°C, K, °F)
Heat flows spontaneously from higher to lower temperature only
K = °C + 273; °F = (9/5)°C + 32
Q = mcΔT — use when temperature changes (no phase change)
Q = mL — use when state changes (temperature stays constant)
Three modes: Conduction (solids, molecular vibration), Convection (fluids, bulk flow), Radiation (no medium, EM waves)
Specific heat of water = 4186 J/kg°C — highest of common substances
Latent heat of fusion of water = 3.34×10⁵ J/kg; vaporisation = 2.26×10⁶ J/kg
Absolute zero = 0 K = −273.15°C — lowest possible temperature

Exam Tips

Always convert mass to kg and temperature to °C (or K) before substituting
Identify WHICH formula to use: Q=mcΔT (temp change) vs Q=mL (phase change)
Draw and label the heating curve — flat regions are phase changes, sloped regions are temperature changes
Name and explain all three modes of heat transfer with one example each
Know why steel feels colder than wood at same temperature (conductivity, not temperature)

Heat Temperature Class 9 Physics CBSE Thermodynamics Conduction Convection Radiation Latent Heat