HeatClass 10 Science Part 1 Chapter 5 Notes – Maharashtra Board

Summary: Heat

Put a glass of ice on the gas and keep a thermometer in it. You might expect the reading to climb steadily, but it sits at 0 °C until every piece of ice has melted. Later it sits at 100 °C while the water boils away. The heat supplied at these flat stages is called latent heat, the hidden heat that breaks bonds between molecules instead of raising the temperature.

Pressure can also change how ice behaves. Press two ice cubes together and they join; hang a loaded wire over an ice slab and it cuts slowly through, yet the slab stays in one piece. Pressure lowers the melting point, the ice melts under the wire, and the water freezes again above it. This melt-and-refreeze trick is called regelation.

Water breaks the usual rule that liquids shrink on cooling. Between 4 °C and 0 °C it actually expands, so water is densest at 4 °C. Hope's apparatus shows this clearly in the lab. In nature, this odd habit lets fish survive under frozen lakes, but it also bursts water pipes and cracks rocks in very cold places.

Air always carries some water vapour. How much it can hold depends on its temperature. When moist air cools to its dew point, droplets appear, which is why a cold bottle from the fridge 'sweats' and grass is wet on winter mornings. Relative humidity tells us how close the air is to being saturated, and explains why Mumbai feels sticky while Nagpur in summer feels dry.

Finally, different materials need different amounts of heat to warm up by the same amount. This property is specific heat capacity. Using Q = m × c × ΔT and the principle of heat exchange (heat lost by the hot body = heat gained by the cold body), we can solve calorimeter problems and even find the specific heat of an unknown metal.

Key points to remember

• During a change of state, a substance absorbs or gives out heat while its temperature stays constant; this heat is latent heat.
• Specific latent heat of fusion of ice = 80 cal/g (333 kJ/kg); specific latent heat of vaporization of water = 540 cal/g (2256 kJ/kg).
• Melting point, boiling point and latent heat of a substance depend on the surrounding pressure.
• Regelation: ice melts under pressure (its melting point drops below 0 °C) and freezes again when the pressure is removed.
• Water contracts when heated from 0 °C to 4 °C; this is the anomalous behaviour of water.
• Water has minimum volume and maximum density at 4 °C.
• Hope's apparatus is used to show that water has maximum density at 4 °C.
• Dew point is the temperature at which unsaturated air, on cooling, becomes saturated with water vapour.
• Absolute humidity = mass of water vapour per unit volume of air (unit kg/m³).
• Relative humidity at the dew point is 100%; above about 60% air feels humid, below 60% it feels dry.
• 1 cal is the heat needed to raise 1 g of water from 14.5 °C to 15.5 °C; 1 kcal = 1000 cal; 1 cal = 4.18 J.
• Heat absorbed or released, Q = m × c × ΔT, where c is the specific heat capacity.
• Water has the highest specific heat in the textbook table (1 cal/g °C); mercury has one of the lowest (0.033 cal/g °C).
• Principle of heat exchange: in an isolated system, heat lost by the hot body = heat gained by the cold body.

Heat – Notes

1. Latent heat — heating ice to steam

• When we heat ice steadily and note the temperature every minute, the graph is not a straight rising line; it has two flat steps.
• First flat step (0 °C): ice and water exist together. Heat goes in, but the temperature does not rise until all ice melts. This temperature is the melting point of ice.
• The absorbed heat is used to loosen the bonds holding molecules in the solid, so it does not show up as a temperature rise. It is 'hidden', hence the name latent heat.
• Next, the water warms from 0 °C to 100 °C. Then a second flat step appears at 100 °C, the boiling point, while water turns into steam.
• Latent heat of fusion: heat absorbed at constant temperature when a solid becomes a liquid. Latent heat of vaporization: heat absorbed at constant temperature when a liquid becomes a gas.
• Specific latent heat means the heat needed for unit mass (1 g or 1 kg). For ice, L_fusion = 80 cal/g; for water, L_vap = 540 cal/g.
• The reverse changes (steam to water, water to ice) give out the same amount of latent heat to the surroundings. That is why a steam burn is far more painful than a burn from boiling water.
• Different substances have different melting points, boiling points and latent heats, and all of these change with atmospheric pressure.

2. Regelation

• Experiment: place an ice slab on a stand and hang equal weights on both ends of a thin metal wire laid across it.
• Observation: the wire slowly sinks through the slab and finally comes out at the bottom, but the slab does not split into two pieces.
• Reason: pressure from the wire lowers the melting point of ice below 0 °C, so the ice just under the wire melts even though it is at 0 °C.
• The water moves above the wire. There the pressure is gone, the melting point returns to 0 °C, and the water freezes again, sealing the cut.
• Regelation: melting of ice under pressure and its refreezing when the pressure is removed.
• Latent heat is involved: heat taken for melting below the wire is given back when the water refreezes above it.
• Everyday example: a gola (ice-ball) vendor presses crushed ice around a stick and it becomes one solid ball; two ice cubes pressed together stick to each other.

3. Anomalous behaviour of water

• Most liquids expand when heated and contract when cooled. Water follows this rule above 4 °C.
• Between 0 °C and 4 °C water does the opposite: on heating from 0 °C it contracts, and on cooling from 4 °C to 0 °C it expands.
• This unusual behaviour of water between 0 °C and 4 °C is called the anomalous behaviour of water.
• The volume–temperature graph of 1 kg water has its lowest point at 4 °C, so water has minimum volume and maximum density at 4 °C.
• Aquatic life: in cold countries, the surface of a lake cools first. Water at 4 °C, being densest, sinks to the bottom. Colder water stays on top and freezes. The ice layer insulates the water below, which stays near 4 °C, so fish and plants survive.
• Damage: water trapped in pipes or in cracks of rocks expands when it cools below 4 °C and freezes. The huge force bursts pipes and splits rocks (helps weathering of rocks in Himalayan regions).

4. Hope's apparatus

• It is a tall metal cylinder with a flat bowl fitted round its middle, and two thermometers: T₁ near the bottom and T₂ near the top.
• The cylinder is filled with water and the bowl is filled with a freezing mixture (ice + salt), which cools the middle part of the water.
• Readings of T₁ and T₂ are noted every 30 seconds and plotted against time.
• At first, the cooled water in the middle becomes denser and sinks, so T₁ (lower) falls quickly while T₂ (upper) falls slowly.
• T₁ stops at 4 °C and stays there. Water cooled below 4 °C becomes lighter and rises, so now T₂ falls fast and reaches 0 °C first.
• Only later does T₁ slowly drop to 0 °C. The point where the two curves cross gives the temperature of maximum density, 4 °C.

5. Dew point and humidity

• About 71% of Earth's surface is water, so evaporation keeps adding water vapour to the air. The moisture in air is called humidity.
• A given volume of air at a given temperature can hold only a limited amount of vapour. Air holding this maximum is saturated; air holding less is unsaturated.
• Warm air can hold more vapour than cold air: at 40 °C about 49 g per kg of dry air, but at 20 °C only about 14.7 g.
• Dew point: the temperature at which unsaturated air, when cooled, becomes saturated. Below this, extra vapour condenses into droplets. The dew point depends on how much vapour the air contains.
• A cold bottle from the fridge gets droplets outside because the air touching it cools below its dew point. Dew on grass and on car windows in the early morning forms the same way.
• Absolute humidity = mass of vapour in unit volume of air (kg/m³). Relative humidity compares the actual vapour to the vapour needed for saturation at that temperature.
• Relative humidity is 100% at the dew point. Above about 60% air feels sticky (Mumbai in monsoon); below 60% it feels dry.
• The white trail behind a high-flying jet is engine vapour that has condensed. In humid air it lasts long; in dry air it is short or does not form at all.

6. Units of heat

• SI unit of heat is the joule (J); CGS unit is the calorie (cal).
• 1 calorie: heat needed to raise the temperature of 1 g of water by 1 °C, from 14.5 °C to 15.5 °C.
• 1 kilocalorie: heat needed to raise 1 kg of water by 1 °C, from 14.5 °C to 15.5 °C. So 1 kcal = 1000 cal.
• The temperature interval is fixed because the heat needed for a 1 °C rise is slightly different in different temperature ranges.
• Relation: 1 cal = 4.18 J. The joule is named after James Prescott Joule, who showed that heat is a form of energy and that energy can change from one form to another.
• Food packets in India show energy in kcal (often written 'Calories'); multiply by 4.18 to get kJ.

7. Specific heat capacity

• Activity: iron, copper and lead balls of equal mass are kept in boiling water so all reach 100 °C, then placed on a thick slab of wax.
• The iron ball sinks deepest, copper goes to a middle depth and lead the least. Iron absorbed and gave out the most heat for the same temperature change.
• So the heat taken by a body depends on its material. This property is specific heat capacity (c).
• Definition: the heat needed to raise the temperature of unit mass of a substance by 1 °C.
• Units: SI — J/kg °C; CGS — cal/g °C. For water c = 1 cal/g °C = 1 kcal/kg °C.
• Textbook values (cal/g °C): water 1.0, paraffin 0.54, kerosene 0.52, aluminium 0.215, iron 0.110, copper 0.095, silver 0.056, mercury 0.033.
• High c of water is why coastal cities like Mumbai have milder climates, and why water is used in car radiators and hot-water bags.
• Heat gained or lost when temperature changes by ΔT: Q = m × c × ΔT.

8. Principle of heat exchange

• When a hot body and a cold body touch, heat flows from the hot body to the cold one.
• The hot body cools and the cold body warms until both reach the same final temperature.
• If both are kept in a heat-resistant (insulated) box, no heat enters or leaves the box; all exchange is between the two bodies.
• Principle of heat exchange: heat lost by the hot body = heat gained by the cold body.
• It is really the law of conservation of energy applied to heat.
• Daily example: mixing hot water from a geyser with cold tap water in a bucket gives warm bathing water at an in-between temperature.

9. Measuring specific heat — mixing method and calorimeter

• A calorimeter is an insulated copper vessel used to measure heat exchanged.
• A hot solid of known mass and temperature is dropped into water of known mass and temperature inside the calorimeter.
• Heat flows until the solid, the water and the calorimeter all reach one common temperature.
• Heat lost by solid (Q) = heat gained by water (Q₁) + heat gained by calorimeter (Q₂).
• Each term is mass × specific heat × temperature change. Knowing c of water and of the calorimeter, the only unknown, c of the solid, can be found.
• Always write 'decrease in temperature' for the hot body and 'increase in temperature' for the cold bodies so that every heat term is positive.

Formulas at a glance

Key terms

Heat – Flashcards

Read the question, answer it in your head, then tap the card to check. 20 cards.

Heat – MCQ Quiz

15 questions. Tap an option — you will see at once if it is right and why.

How to study Heat

⏱ About 4 hours over 4 days

1. Day 1 — Latent heat and regelation
   Read the latent heat and regelation sections. Draw the temperature–time graph for ice → water → steam and label each part (ice + water at 0 °C, water warming, boiling at 100 °C). Write the definitions of both specific latent heats with values 80 and 540 cal/g. Explain regelation aloud using the wire-and-ice set-up.
2. Day 2 — Anomalous behaviour and Hope's apparatus
   Draw the volume–temperature graph of water with its minimum at 4 °C. Draw and label Hope's apparatus, then the T₁/T₂ time graph. Write 3-mark explanations for aquatic life in frozen lakes and bursting pipes / cracking rocks.
3. Day 3 — Humidity, dew point and units
   Learn saturated vs unsaturated air, dew point, absolute and relative humidity. Relate dew formation to droplets on a cold bottle. Memorise the calorie definition with its 14.5–15.5 °C range and 1 cal = 4.18 J. Do our flashcards for definitions.
4. Day 4 — Specific heat and numericals
   Understand the wax-slab activity and Q = m c ΔT. Solve every solved example in these notes on paper without looking, including mixed latent + specific heat problems and a calorimeter problem. End with our 15-question quiz and revise anything you got wrong.

✍️ Diagrams to practise

• Temperature vs time graph for heating ice to steam (with flat parts at 0 °C and 100 °C)
• Regelation experiment — ice slab with wire and two weights
• Volume vs temperature graph of water (minimum at 4 °C)
• Hope's apparatus with thermometers T₁ and T₂ and freezing mixture
• Time vs temperature graph from Hope's apparatus (T₁ and T₂ curves)
• Specific heat activity — iron, copper and lead balls on a wax slab