Heat — Class 10 Physical Science
"Heat is ENERGY in transit. Temperature is the MEASURE of that energy. They are not the same — and confusing them is the most common mistake in thermodynamics."
1. Heat vs Temperature
| Heat | Temperature |
|---|---|
| ENERGY transferred due to temperature difference | MEASURE of hotness or coldness |
| Unit: JOULE (J) | Unit: Celsius (°C), Kelvin (K) |
| Depends on MASS | Does NOT depend on mass |
| A SPOON of water at 100°C has MORE temperature but LESS heat than a BATHTUB of water at 40°C | — |
'Heat flows SPONTANEOUSLY from HOT to COLD. Never the reverse (Second Law of Thermodynamics).'
2. Specific Heat Capacity (c)
Definition: The amount of HEAT required to raise the temperature of 1 kg of a substance by 1°C (or 1 K). Unit: J/kg°C or J/kgK.
Formula: Q = mcΔθ Q = heat energy (J). m = mass (kg). c = specific heat capacity. Δθ = change in temperature.
Why Water Is the BEST Coolant
Water has one of the HIGHEST specific heat capacities: 4200 J/kg°C. 'Water can absorb a LOT of heat without its temperature rising much. This is why: car radiators use water. Coastal areas have MODERATE temperatures (the sea heats and cools SLOWLY). Your body (65% water) resists temperature changes.'
Comparing Substances
| Substance | c (J/kg°C) |
|---|---|
| Water | 4200 |
| Iron | 450 |
| Aluminium | 900 |
| Sand | 800 |
'Sand heats up FASTER than water (lower c). That's why: deserts are HOT during the day and COLD at night — sand heats and cools quickly. The beach sand burns your feet — but the sea is pleasantly cool.'
3. Latent Heat
Definition: Heat absorbed or released during a CHANGE OF STATE — WITHOUT any change in TEMPERATURE. 'During a phase change: ALL the heat energy goes into breaking INTERMOLECULAR BONDS — NOT into raising temperature. This is why ice water stays at 0°C until ALL the ice has melted — no matter how much you heat it.'
| Type | Change | Value (for water) |
|---|---|---|
| Latent Heat of Fusion (L_f) | Solid ↔ Liquid | 334 J/g (or 3.34 × 10⁵ J/kg) |
| Latent Heat of Vaporisation (L_v) | Liquid ↔ Gas | 2260 J/g (or 2.26 × 10⁶ J/kg) |
Total heat for ice (−t°C) → steam (100°C): Q = mcᵢΔθ₁ + mL_f + mc_wΔθ₂ + mL_v. 'The VAST majority of energy goes into vaporisation — turning 100°C water into 100°C steam takes MORE ENERGY than heating that water from 0°C to 100°C. Steam burns are MORE SEVERE than boiling water burns — because steam carries latent heat.'
4. Calorimetry Principle
Heat LOST by hot body = Heat GAINED by cold body + calorimeter (assuming no heat loss to surroundings). Used to find: Specific heat of unknown substances. Final temperature of mixtures.
5. Evaporation — The Cooling Effect
Evaporation is a SURFACE phenomenon — occurs at ALL temperatures (unlike boiling, which occurs at a SPECIFIC temperature throughout the liquid). Molecules with HIGH kinetic energy ESCAPE the surface → average KE of remaining molecules DECREASES → temperature DROPS. 'This is why: sweating cools you. Water in an earthen pot (matka) stays cool — water seeps out, evaporates, cools the pot. You feel cold after a shower — water evaporating from your skin.'
Factors Affecting Evaporation Rate
Temperature (↑ → faster). Surface Area (↑ → faster — clothes spread out dry faster). Humidity (↑ → slower — already saturated air cannot absorb more). Wind Speed (↑ → faster — wind removes water vapour, maintaining gradient).
6. Common Mistakes
- 'Heat and temperature are the same' — Heat is ENERGY. Temperature is its MEASURE. An iceberg has MORE heat energy than a cup of boiling water — because of its MASS.
- 'Temperature changes during a phase change' — It does NOT. At 0°C, ice-water mixture stays at 0°C until ALL ice melts. The energy goes into breaking bonds, not raising temperature.
- 'Evaporation = boiling' — Evaporation: SURFACE, any temperature. Boiling: BULK, specific temperature (100°C for water at 1 atm).
7. AP SSC Exam Focus
| Topic | Marks |
|---|---|
| Q = mcΔθ problems | 3-4 |
| Latent heat concept | 2-3 |
| Evaporation factors | 2-3 |
| Calorimetry | 3-4 |
8. Worked Numerical Problems
Specific Heat Capacity Numericals
Example 1: How much heat is required to raise the temperature of 2 kg of water from 20°C to 50°C? (c_w = 4200 J/kg°C) Solution: Q = mcΔθ = 2 × 4200 × (50 − 20) = 2 × 4200 × 30 = 252,000 J = 252 kJ.
Example 2: A 0.5 kg iron pan (c = 450 J/kg°C) is heated from 25°C to 200°C. Find the heat absorbed. Solution: Q = 0.5 × 450 × (200 − 25) = 0.5 × 450 × 175 = 39,375 J ≈ 39.4 kJ.
Example 3: 200 g of a metal at 100°C is placed in 100 g of water at 30°C. The final temperature is 40°C. Find the specific heat capacity of the metal. (c_w = 4200 J/kg°C) Solution: Heat LOST by metal = Heat GAINED by water. m_m × c_m × Δθ_m = m_w × c_w × Δθ_w. 0.2 × c_m × (100 − 40) = 0.1 × 4200 × (40 − 30). 0.2 × c_m × 60 = 0.1 × 4200 × 10. 12 × c_m = 4200. c_m = 4200/12 = 350 J/kg°C.
Latent Heat Numericals
Example 4: How much heat is needed to melt 500 g of ice at 0°C? (L_f = 334 J/g) Solution: Q = mL_f = 500 × 334 = 167,000 J = 167 kJ.
Example 5: Calculate the heat required to convert 100 g of ice at −10°C to steam at 100°C. (c_ice = 2100 J/kg°C, L_f = 3.34 × 10⁵ J/kg, c_w = 4200 J/kg°C, L_v = 2.26 × 10⁶ J/kg) Solution: m = 0.1 kg. Step 1: Heat ice from −10°C to 0°C: Q₁ = 0.1 × 2100 × 10 = 2100 J. Step 2: Melt ice at 0°C: Q₂ = 0.1 × 3.34 × 10⁵ = 33,400 J. Step 3: Heat water from 0°C to 100°C: Q₃ = 0.1 × 4200 × 100 = 42,000 J. Step 4: Vaporise water at 100°C: Q₄ = 0.1 × 2.26 × 10⁶ = 226,000 J. Total Q = 2100 + 33,400 + 42,000 + 226,000 = 303,500 J ≈ 303.5 kJ.
9. Methods of Heat Transfer — Detailed
Conduction
Transfer of heat through a substance WITHOUT the movement of the substance itself. Heat travels by VIBRATIONS of atoms/molecules. 'Metals are GOOD conductors of heat (free electrons carry thermal energy). Wood and plastic are POOR conductors (INSULATORS).'
Conduction Formula — Fourier's Law: Q/t = kA(ΔT/L), where Q/t = rate of heat transfer, k = thermal conductivity, A = area, ΔT = temperature difference, L = thickness.
Examples: Handle of a metal spoon gets hot when placed in hot soup. 'Touching a metal chair feels COLDER than a wooden one — even at the same temperature. Metal CONDUCTS heat away from your hand faster.'
Convection
Transfer of heat by the ACTUAL MOVEMENT of a FLUID (liquid or gas). 'Hot fluid expands → becomes LESS dense → RISES. Cold fluid CONTRACTS → becomes MORE dense → SINKS. This creates a CONVECTION CURRENT.'
Examples: Sea breezes (land heats faster → air rises → cool sea air moves inward). Room heaters placed near the floor (hot air rises). Air conditioners placed near the ceiling (cool air sinks). Boiling water — water rises from the bottom, circulates.
Radiation
Transfer of heat WITHOUT any medium. Travels as ELECTROMAGNETIC WAVES (infrared). Can travel through VACUUM. 'Heat from the Sun reaches us through SPACE — 150 million km of VACUUM. This is ONLY possible by RADIATION.'
Properties: All objects EMIT radiation. HOTTER objects emit MORE radiation. DARK, rough surfaces are GOOD absorbers and GOOD emitters. SHINY, white surfaces are POOR absorbers and POOR emitters (good REFLECTORS).
Examples: Solar water heaters (black surface absorbs maximum radiation). White clothes in summer (reflect radiation — keep you cool). Space blankets (shiny surface reflects body heat back).
10. Applications of Heat Transfer
Vacuum Flask (Thermos) : DOUBLE-WALLED glass bottle with VACUUM between walls. Vacuum PREVENTS conduction and convection. SILVERED surfaces reflect radiation (prevent heat loss by radiation). CORK or PLASTIC stopper prevents evaporation and convection at the top. 'A thermos keeps hot things HOT and cold things COLD — it MINIMISES ALL THREE modes of heat transfer.'
Greenhouse Effect: Sun's short-wave radiation passes through GLASS → heats the interior. The interior radiates LONG-WAVE (infrared) radiation → glass TRAPS it (doesn't let it escape) → interior WARM UP. 'The same principle is WARMING the Earth — CO₂ and other greenhouse gases act like the glass of a greenhouse.'
11. Self-Test
Q1: Why does a desert get very hot during the day and very cold at night? A1: Sand has LOW specific heat capacity (~800 J/kg°C). It heats up QUICKLY during the day (high temperature) and cools down QUICKLY at night (low temperature). Water has HIGH specific heat capacity — it resists temperature changes.
Q2: A 2 kg iron block is heated from 30°C to 80°C. If the specific heat of iron is 450 J/kg°C, calculate the heat supplied. A2: Q = mcΔθ = 2 × 450 × (80 − 30) = 2 × 450 × 50 = 45,000 J = 45 kJ.
Q3: Why is steam at 100°C more dangerous than boiling water at 100°C? A3: Steam carries ADDITIONAL LATENT HEAT of vaporisation (2.26 × 10⁶ J/kg). When steam condenses on your skin, it releases this latent heat — causing MORE SEVERE burns than boiling water alone.
Q4: Explain the sea breeze phenomenon. A4: During the day, land heats up FASTER than the sea. Air above the land RISES (convection). Cooler, denser air from the sea MOVES IN to replace it → SEA BREEZE. At night, land cools FASTER → reverse → LAND BREEZE.
Q5: Why do we wear white or light-coloured clothes in summer? A5: White/light colours REFLECT most of the incident solar radiation (poor absorbers). Dark colours ABSORB more radiation → get hotter. 'This is a practical application of RADIATION principles.'
Q6: Ice at 0°C is more effective for cooling than water at 0°C. Why? A6: Ice at 0°C must ABSORB latent heat of fusion (334 J/g) to MELT into water at 0°C. This EXTRA heat absorption makes ice a better cooling agent than the same mass of water at the same temperature.
Q7: In calorimetry, why must we minimise heat loss to the surroundings? A7: The calorimetry principle assumes HEAT LOST = HEAT GAINED. If heat escapes to surroundings, the calculation becomes INACCURATE. To minimise: use a STIRRER (uniform temperature), a LAGGED (insulated) container, and a LID (prevent convection losses).
