Heat Class 10 Notes, Formulas, Questions | SSC

This chapter explains several everyday effects of heat. It shows how a substance absorbs heat without a rise in temperature while changing state (latent heat), how pressure lowers the melting point of ice (regelation), and why water behaves unusually between 0 and 4 degrees C. It then deals with water vapour in air, dew point, absolute and relative humidity, the units of heat, specific heat capacity and the principle of heat exchange used in calorimetry.

Summary

When ice is heated, its temperature stays at 0 degrees C until all of it melts, then the water warms to 100 degrees C, where the temperature again stays constant until all the water turns into steam. The heat absorbed during a change of state at constant temperature is used to weaken or break the bonds between molecules and is called latent heat. The constant temperature at which a solid melts is its melting point, and at which a liquid boils is its boiling point. Heat absorbed per unit mass while a solid turns into liquid is the specific latent heat of fusion (for ice 80 cal/g or 333 kJ/kg), and while a liquid turns into gas is the specific latent heat of vaporisation (for water 540 cal/g or 2256 kJ/kg). Melting point, boiling point and latent heat depend on atmospheric pressure. The same latent heat is given out when gas turns into liquid or liquid into solid.

When a wire with weights at its ends is placed across an ice slab, it slowly passes through the ice and comes out at the bottom, but the slab does not break. Pressure lowers the melting point of ice below 0 degrees C, so the ice under the wire melts; once the wire passes and the pressure is removed, the melting point returns to 0 degrees C and the water freezes again. This melting under pressure and refreezing when pressure is removed is called regelation. It explains how shredded ice is pressed into an ice ball and why two pieces of ice pressed together stick.

Most liquids expand when heated, but water contracts when heated from 0 degrees C to 4 degrees C and expands above 4 degrees C. This is the anomalous behaviour of water. Water therefore has its minimum volume and maximum density at 4 degrees C. Hope's apparatus demonstrates this: a cylinder of water has a freezing mixture of ice and salt around its middle, with thermometers T1 below and T2 above. Cooled water sinks, so T1 falls quickly to 4 degrees C and stays there; water cooled below 4 degrees C becomes lighter and rises, so T2 then falls to 0 degrees C first. The point where the two temperature curves cross shows the temperature of maximum density.

Because of this behaviour, in cold regions the top layer of a lake freezes while the water below stays at about 4 degrees C, so aquatic plants and animals survive. Water expands on cooling below 4 degrees C and on freezing, so water pipes burst and rocks crack in cold winters.

About 71% of the earth's surface is water, so the air always contains water vapour; this moisture is humidity. At a given temperature air can hold only a limited amount of vapour; air holding the maximum is saturated, and air holding less is unsaturated. Warm air can hold more vapour (49 g per kg of dry air at 40 degrees C against 14.7 g at 20 degrees C). The temperature at which unsaturated air becomes saturated on cooling is the dew point. Absolute humidity is the mass of vapour per unit volume of air (kg/m^3). Relative humidity is the ratio of the actual mass of vapour in a volume of air to the mass needed to saturate it at that temperature, usually expressed as a percentage. It is 100% at the dew point; above 60% the air feels humid and below 60% it feels dry. Droplets on a cold bottle from the refrigerator, dew on grass and the white trail of an aeroplane are due to condensation of vapour.

The SI unit of heat is the joule and the CGS unit is the calorie. One calorie is the heat needed to raise the temperature of 1 g of water by 1 degree C from 14.5 degrees C to 15.5 degrees C; one kcal does the same for 1 kg, so 1 kcal = 1000 cal and 1 cal = 4.18 J. A definite temperature interval is chosen because the heat needed differs slightly in other ranges. The unit joule is named after James Prescott Joule, who showed that energy can be converted from one form to another.

When iron, copper and lead spheres of equal mass heated to 100 degrees C are placed on a wax slab, the iron sphere sinks deepest and lead the least, showing that equal masses absorb different amounts of heat for the same temperature rise. Specific heat capacity is the heat needed to raise the temperature of unit mass of a substance by 1 degree C. Its SI unit is J/kg degree C and CGS unit cal/g degree C; water has the high value 1 cal/g degree C. Heat gained or lost is Q = m c delta T.

When a hot and a cold object exchange heat in an insulated system, heat flows from the hot to the cold body until both reach the same temperature. The principle of heat exchange states that heat lost by the hot object equals heat gained by the cold object. In the mixing method, a hot solid is dropped into water in a calorimeter; heat lost by the solid equals heat gained by the water plus heat gained by the calorimeter, from which the specific heat of the solid can be found.

Mind map

Heating ice from 0 degrees C to steam (temperature versus time graph)

Important definitions

Formulas

Key points to remember

• During a change of state temperature stays constant though heat is absorbed or released.
• Ice melts at 0 degrees C and water boils at 100 degrees C at normal pressure; these values depend on pressure.
• Latent heat of fusion of ice is 80 cal/g (333 kJ/kg); latent heat of vaporisation of water is 540 cal/g (2256 kJ/kg).
• Latent heat is given out when vapour condenses or liquid freezes.
• Pressure lowers the melting point of ice; this causes regelation.
• Water has minimum volume and maximum density at 4 degrees C.
• In Hope's apparatus the lower thermometer stays at 4 degrees C; the upper one reaches 0 degrees C first.
• Relative humidity above 60% feels humid, below 60% feels dry; it is 100% at dew point.
• Specific heat of water (1 cal/g degree C) is higher than that of metals like iron (0.110), copper (0.095) and mercury (0.033).
• In the wax slab experiment the iron sphere sinks deepest, then copper, then lead.
• Specific heat is measured by the mixing method using a calorimeter.

Solved examples

Example 1. How much heat is needed to raise the temperature of 5 kg of water from 20 degrees C to 100 degrees C?

\( m = 5 \) kg, \( c = 1 \) kcal/kg degree C, \( \Delta T = 100 - 20 = 80 \) degree C. \( Q = mc\Delta T = 5\times1\times80 = 400 \) kcal.

Example 2. A 100 g copper sphere at 100 degrees C is dropped into 195 g of water at 20 degrees C in a 50 g copper calorimeter. Specific heat of copper = 0.1 cal/g degree C. Find the final temperature.

Let final temperature be T. Heat lost by sphere = heat gained by water + calorimeter: \( 100\times0.1\times(100 - T) = 195\times1\times(T - 20) + 50\times0.1\times(T - 20) \). So \( 10(100 - T) = 200(T - 20) \Rightarrow 1000 - 10T = 200T - 4000 \Rightarrow 210T = 5000 \Rightarrow T \approx 23.8 \) degree C.

Example 3. 80 g of steam at 97 degrees C is released on an ice slab at 0 degrees C. How much ice melts? (L ice = 80 cal/g, L vap = 540 cal/g)

Heat released by steam condensing: \( 80\times540 \) cal. Heat released by water cooling from 97 to 0 degree C: \( 80\times1\times97 \) cal. Total \( = 80(540 + 97) = 80\times637 = 50960 \) cal. Ice melted \( m = \dfrac{50960}{80} = 637 \) g.

Example 4. Equal heat is given to two objects A and B of mass 1 g. A's temperature rises by 3 degrees C and B's by 5 degrees C. Which has more specific heat and by what factor?

\( Q = mc\Delta T \) with equal Q and m, so \( c \propto \dfrac{1}{\Delta T} \). \( \dfrac{c_A}{c_B} = \dfrac{5}{3} \). A has the greater specific heat, 5/3 times that of B.

Example 5. Water at 20 degrees C is to be made into 2 kg of ice at 0 degrees C using ammonia (latent heat of vaporisation 341 cal/g). How much ammonia must evaporate?

Heat to remove: cooling \( 2000\times1\times20 = 40000 \) cal, freezing \( 2000\times80 = 160000 \) cal; total 200000 cal. Ammonia needed \( = \dfrac{200000}{341} \approx 586.5 \) g (about 586.4 g as in the textbook).

Example 6. A pot has 150 g of ice at 0 degrees C. How much steam at 100 degrees C must be mixed so that the final water is at 50 degrees C?

Heat gained by ice: melting \( 150\times80 = 12000 \) cal, warming \( 150\times1\times50 = 7500 \) cal; total 19500 cal. Heat lost by steam of mass m: condensing \( 540m \), cooling to 50 degree C \( 50m \); total \( 590m \). \( 590m = 19500 \Rightarrow m \approx 33 \) g.

Example 7. A 100 g calorimeter (c = 0.1 kcal/kg degree C) contains 250 g of liquid (c = 0.4 kcal/kg degree C) at 30 degrees C. 10 g of ice at 0 degrees C is added. Find the final temperature.

In cal/g degree C, calorimeter c = 0.1, liquid c = 0.4. Heat lost: \( (250\times0.4 + 100\times0.1)(30 - T) = 110(30 - T) \). Heat gained by ice: \( 10\times80 + 10\times1\times T = 800 + 10T \). \( 3300 - 110T = 800 + 10T \Rightarrow 120T = 2500 \Rightarrow T \approx 20.8 \) degree C.

Important questions and answers

Q1. What is latent heat? How does the state change if latent heat is given off?

Latent heat is the heat absorbed or given out at constant temperature during a change of state; it is used to break or form intermolecular bonds. When latent heat is given off, a gas changes into liquid (condensation) or a liquid changes into solid (freezing).

Q2. Explain regelation with an experiment.

Place an ice slab on a stand and hang a metal wire with equal weights at both ends across it. The wire slowly passes through the ice, but the slab does not break. The pressure under the wire lowers the melting point of ice below 0 degrees C, so the ice there melts. Above the wire the pressure is removed, the melting point returns to 0 degrees C and the water freezes again. This is regelation.

Q3. What is the anomalous behaviour of water? How is it studied using Hope's apparatus?

Water contracts on heating from 0 to 4 degrees C instead of expanding, so its density is maximum at 4 degrees C. In Hope's apparatus a freezing mixture surrounds the middle of a cylinder of water, with thermometers T1 at the bottom and T2 at the top. Cooled water, being denser, sinks, so T1 falls quickly to 4 degrees C and stays there. Water cooled below 4 degrees C becomes lighter and rises, so T2 falls to 0 degrees C first, and T1 reaches 0 degrees C later. The crossing point of the two curves gives 4 degrees C, the temperature of maximum density.

Q4. Explain the role of the anomalous behaviour of water in preserving aquatic life in cold regions.

As the air cools, the surface water cools and sinks until the whole water is at 4 degrees C. Water cooled below 4 degrees C is lighter and stays at the top, where it freezes. Ice floats and, being a poor conductor of heat, insulates the water below, which remains at about 4 degrees C, so aquatic plants and animals survive.

Q5. Give reason: In cold regions in winter, water pipes burst and rocks crack.

Water in pipes and in the cracks of rocks expands when cooled below 4 degrees C and expands further on freezing. This expansion exerts great pressure on the pipes and rocks, so the pipes burst and rocks crack.

Q6. Relate the water droplets on a bottle taken out of the refrigerator with the formation of dew.

The cold bottle cools the air near it below its dew point, so the air becomes saturated and the excess vapour condenses into droplets on the bottle. Similarly, in the early morning the air near grass, leaves and glass cools below the dew point and the vapour condenses as dew.

Q7. What is specific heat capacity? How will you show that different substances have different specific heat capacities?

It is the heat required to raise the temperature of unit mass of a substance by 1 degree C. Heat equal mass spheres of iron, copper and lead in boiling water to 100 degrees C and place them on a thick wax slab. The iron sphere sinks deepest, copper less and lead the least. Each gave out different heat while cooling through the same range, so their specific heat capacities differ (iron highest, lead lowest).

Q8. While defining the unit of heat, which temperature interval is chosen, and why?

The interval 14.5 degrees C to 15.5 degrees C is chosen. The heat needed to raise the temperature of 1 g of water by 1 degree C differs slightly in different ranges, so a definite range must be specified.

Q9. State the principle of heat exchange. Which property is measured using it?

When hot and cold objects exchange heat in an isolated system, heat lost by the hot object equals heat gained by the cold object; heat flows from the hot to the cold body until both reach the same temperature. It is used to measure the specific heat capacity of a substance by the mixing method using a calorimeter.

Q10. Define dew point and relative humidity. How will you decide whether air is saturated?

Dew point is the temperature at which unsaturated air becomes saturated on cooling. Relative humidity is the ratio of the actual mass of vapour in a volume of air to the mass needed to saturate it at that temperature, times 100. If the relative humidity is 100%, the air is saturated; at the dew point it is 100%. If it is less than 100%, the air is unsaturated.

Q11. Explain the temperature versus time graph obtained on heating ice.

Part AB is horizontal at 0 degrees C: ice melts and absorbs latent heat of fusion without a rise in temperature. Part BC rises from 0 to 100 degrees C as the water warms. Then the graph is again horizontal at 100 degrees C: water boils and absorbs latent heat of vaporisation at constant temperature.

Exam tips

• Keep units consistent: use cal, g and degree C together, or J, kg and degree C together.
• In mixing problems write 'heat lost = heat gained' first, then list every body (solid, water, calorimeter, ice melting, melted ice warming).
• When ice is involved, remember both steps: melting (mL) and warming of the melted water (mc delta T).
• Draw and label the ice to steam graph and the Hope's apparatus graph; they are frequently asked.
• Memorise: L fusion ice 80 cal/g, L vaporisation water 540 cal/g, c water 1 cal/g degree C, 1 cal = 4.18 J.

FAQs

Why does temperature not rise while ice melts?

The heat absorbed is used to weaken the bonds between water molecules to change the solid into liquid, not to raise the temperature. This heat is latent heat.

At what temperature is the density of water maximum?

At 4 degrees C, where its volume is minimum.

When does air feel humid?

When its relative humidity is more than 60%; below 60% the air feels dry.

Why does an aeroplane sometimes leave a long white trail in winter?

Vapour from the engine condenses into a cloud-like trail. If the surrounding air has high relative humidity, the trail takes a long time to disappear; in drier air it is small or does not form.

These notes were prepared with AI help from the Balbharati Science and Technology Part 1 textbook (2026-27) and checked against it.