Calorimetry Class 10 ICSE — Selina Chapter 11

Calorimetry Class 10 ICSE (Selina Chapter 11) explains how heat energy is measured, stored and transferred between bodies. The chapter moves from the meaning of heat and temperature through heat capacity and specific heat capacity, then into phase changes and latent heat. Every definition carries its formula and a worked numerical pattern you can apply to the exercises in this ICSE chapter. The goal is simple: you finish knowing which equation to use for each problem.

Download Calorimetry Class 10 ICSE Physics Chapter 11 PDF

Download the Calorimetry Class 10 ICSE Physics chapter PDF from the Selina Concise textbook to revise offline or print for practice. It is the full chapter as printed in the book — 22 pages, 26 sections, 59 questions.

What is inside Chapter 11

22 pages · 26 sections · 2 worked examples · 59 exercise questions

• EXERCISE-11(A): 29 questions
• MULTIPLE CHOICE TYPE: 16 questions
• MULTIPLE CHOICE TYPE: 14 questions

11.1 CONCEPT OF HEAT

Heat is the internal energy of the molecules that make up a body. Molecules in any substance move randomly, so they hold kinetic energy. They also pull on each other, so they hold potential energy. The sum of both is the internal energy. Heat is this internal energy flowing from a hot body to a cold body when they touch.

Calorimetry is the measurement of the quantity of heat. The SI unit is the joule (J). One calorie (cal) raises 1 g of water through 1°C. The precise 15°C calorie uses 14.5°C to 15.5°C. The conversion is 1 cal = 4.186 J, usually taken as 4.2 J. One kilocalorie (kcal) equals 1000 cal, about 4200 J, and is used for food energy values.

Worked pattern: heating 200 g of water from 20°C to 21°C needs 200 cal. Converting to joules gives 200 \times 4.2 = 840\ \text{J}. This is the same route every Part A numerical takes.

11.2 CONCEPT OF TEMPERATURE

Temperature tells the thermal state of a body, its degree of hotness or coldness. It decides the direction of heat flow. When two bodies at different temperatures touch, heat moves from the higher to the lower. Equal temperature means no net heat transfer, but it does not mean equal heat content. Temperature alone does not reveal how much heat a body holds.

The SI unit is the kelvin (K). The relation is T\ \text{K} = 273 + t\ ^\circ\text{C}, or more precisely 273.15. Absolute zero is 0 K, about -273.15°C, where molecular motion ceases. A temperature difference is numerically the same on both scales: \Delta t\ ^\circ\text{C} = \Delta T\ \text{K}.

Common mistake: adding 273 to a temperature difference. A rise of 15°C is a rise of 15 K. The conversion applies to a single reading, not to a change. Also, receiving heat raises temperature only when the physical state and dimensions stay unchanged.

11.3 FACTORS AFFECTING THE QUANTITY OF HEAT ABSORBED TO INCREASE THE TEMPERATURE OF A BODY

Three factors decide the heat absorbed to raise a body’s temperature: the mass, the rise in temperature, and the material. Experiments show heat is directly proportional to mass, Q \propto m, and directly proportional to the rise, Q \propto \Delta t. Combining them gives the central equation:

Q = cm\Delta t

Here c is the specific heat capacity, a characteristic of the substance. Equal masses of water and copper heated through the same rise show why: water absorbs nearly ten times the heat of copper. In the Class 10 ICSE syllabus, this equation powers most Part A numericals.

Worked pattern: heat needed to raise 2.5 kg of water from 20°C to 45°C. The rise is \Delta t = 45 - 20 = 25\ ^\circ\text{C} = 25\ \text{K}. Then:

Q = 2.5 \times 4200 \times 25 = 262500\ \text{J}

The units cancel to joules: \text{kg} \times \text{J kg}^{-1}\text{K}^{-1} \times \text{K} = \text{J}.

11.4 DIFFERENCE BETWEEN HEAT AND TEMPERATURE

This section gives the six-point comparison that answers Exercise 11(A) Question 6.

The key idea: heat is an amount, temperature is a state. A small hot object and a large lukewarm object can hold the same heat but sit at very different temperatures.

11.5 THERMAL (OR HEAT) CAPACITY (C’ = Q / Delta T)

Heat capacity is the heat energy required to raise the whole body’s temperature by 1°C (or 1 K). It is denoted C'. The formula is:

C' = \frac{Q}{\Delta t}

The SI unit is J K⁻¹. Common units are cal °C⁻¹ and kcal °C⁻¹. The relations are 1 kcal °C⁻¹ = 1000 cal °C⁻¹ and 1 cal K⁻¹ = 4.2 J K⁻¹. A vessel with heat capacity 30 J K⁻¹ needs 30 J for each kelvin rise.

Worked pattern: a pan of heat capacity 240 J K⁻¹ receives 1200 J. Its temperature rises by \Delta t = 1200 / 240 = 5\ \text{K}. Heat capacity depends on both the substance and the mass of the body.

11.6 SPECIFIC HEAT CAPACITY

Specific heat capacity is heat capacity per unit mass. It is denoted c. The definitions give:

c = \frac{C'}{m} = \frac{Q}{m\Delta t}

In words: specific heat capacity is the heat energy to raise unit mass of the substance by 1°C (or 1 K). The working relation is Q = mc\Delta t, and the link C' = mc. The SI unit is J kg⁻¹ K⁻¹.

Unit conversions: 1 cal g⁻¹ °C⁻¹ = 1 kcal kg⁻¹ K⁻¹ = 4.2 × 10³ J kg⁻¹ K⁻¹. For copper, c = 0.4\ \text{J g}^{-1}\text{K}^{-1} means 1 g of copper needs 0.4 J per kelvin.

Common mistake: writing mass in grams while using c in J kg⁻¹ K⁻¹. Convert grams to kilograms first. For example, 50 g becomes 0.05 kg before substitution.

11.7 DISTINCTION BETWEEN THE HEAT CAPACITY AND SPECIFIC HEAT CAPACITY

Quick check: if the size or mass of the body matters, it is heat capacity. If only the material matters, it is specific heat capacity. This answers Exercise 11(A) Question 11, which asks for three differences.

11.8 SPECIFIC HEAT CAPACITY OF SOME COMMON SUBSTANCES

Specific heat capacity is a characteristic property. A good conductor has low specific heat capacity; a bad conductor has high specific heat capacity. When equal masses of two substances are heated on the same burner, the one with lower c shows a faster, higher rise in temperature.

Water is unusual: c = 4200\ \text{J kg}^{-1}\text{K}^{-1}. The same substance differs by phase: water 4200, ice 2100, steam 460 J kg⁻¹ K⁻¹. Hydrogen has the maximum value at 14630 J kg⁻¹ K⁻¹.

11.9 CALORIMETER

A calorimeter is a cylindrical vessel used to measure heat gained or lost by a body when mixed with another. It is made of a thin sheet of copper for two reasons. Copper conducts heat well, so the vessel quickly takes the temperature of its contents. Copper also has low specific heat, so the calorimeter steals very little heat.

Each design part reduces a specific heat loss. The polished surfaces reduce radiation. The wooden jacket and wool or cotton packing reduce conduction. The wooden lid reduces convection. The stirrer mixes the contents evenly, and the thermometer hole lets you read the temperature.

11.9 PRINCIPLE OF METHOD OF MIXTURES (OR PRINCIPLE OF CALORIMETRY)

When a hot body mixes with a cold body, heat passes until both reach the same temperature. With perfect insulation, heat lost by the hot body equals heat gained by the cold body. This is the principle of method of mixtures, also called the principle of calorimetry, based on the law of conservation of energy.

For substance A of mass m_1, specific heat c_1, at temperature t_1, mixed with substance B of mass m_2, specific heat c_2, at lower temperature t_2, with final temperature t:

m_1c_1(t_1 - t) = m_2c_2(t - t_2)

If a calorimeter of mass M and specific heat c is used, it also gains heat:

m_1c_1(t_1 - t) = (m_2c_2 + Mc)(t - t_2)

Worked pattern: 500 g of water at 90°C mixed with 300 g of water at 30°C in a vessel of negligible heat capacity. Set 500(90 - t) = 300(t - 30). Solving gives t = 67.5\ ^\circ\text{C}.

Common mistake: forgetting the Mc term for the calorimeter or vessel when it is present. The calorimeter absorbs heat too, and leaving it out changes the answer.

11.10 NATURAL PHENOMENA AND CONSEQUENCES OF HIGH SPECIFIC HEAT CAPACITY OF WATER

Six consequences follow from water’s high specific heat capacity of 4200 J kg⁻¹ K⁻¹.

1. Moderate climate near the sea shore. Sand heats and cools about five times faster than water, so land and sea breezes form.
2. Hot water bottles for fomentation. Water cools slowly because of its high specific heat.
3. Water as an effective coolant. Radiators in cars and generators use water because it extracts heat without rising much in temperature.
4. Water as a heat reservoir. Wine and juice bottles stay unfrozen because water imparts heat before reaching its freezing point.
5. Farmers flood fields on cold nights. The water keeps plant capillaries from freezing and bursting.
6. Plant and animal bodies stay at steady temperature because they contain 80 to 90 percent water.

11.11 SOME EXAMPLES OF HIGH AND LOW HEAT CAPACITY

Three design cases show how heat capacity is used in practice.

• A thick cooking-pan base gives high heat capacity. The pan heats slowly, cooks food properly, and keeps it warm after cooking.
• A thick heavy electric-iron base stays hot long after the current is switched off.
• A thin copper calorimeter has low mass and hence low heat capacity. It takes negligible heat from its contents.

The worked examples in this section of the book show heat capacity and specific heat capacity calculations from given Q, m and Δt. This section answers Exercise 11(A) Question 29.

11.12 CHANGE OF PHASE (STATE)

Change of phase is the change from one state to another at a constant temperature, brought about by exchange of heat. Ice becomes water on heating, and water becomes steam on further heating.

• Melting: solid to liquid
• Freezing: liquid to solid
• Vaporisation: liquid to vapour
• Condensation (or liquefaction): vapour to liquid
• Sublimation: solid to vapour
• Solidification: vapour to solid

11.13 MELTING AND FUSION

Melting is the change from solid to liquid by absorption of heat at a constant temperature. The melting point is that constant temperature. Freezing (or fusion) is the reverse: liquid to solid with rejection of heat. The freezing point is its constant temperature.

Two notes to remember. For a pure substance, the melting point and freezing point are identical. For a given mass, the heat absorbed during melting equals the heat rejected during freezing.

11.14 HEATING CURVE OF ICE DURING MELTING

The experiment: ice chips fill a boiling tube half full. A thermometer sits in the ice with its bulb clear of the walls. Heat the tube slowly and note the temperature every half minute until the melted water reaches 30°C.

Plot temperature on the Y-axis against time on the X-axis. The graph has two parts. Part AB is horizontal at 0°C while the ice melts. Part BC rises as the water warms. During AB, the supplied heat is used in melting the ice, so the temperature stays fixed at the melting point.

11.15 CHANGE IN VOLUME ON MELTING

Most substances like lead and wax expand on melting. Some substances like ice contract. One gram of solid wax at 64°C has volume 1.161 cm³; when melted it becomes 1.166 cm³. One gram of ice at 0°C has volume 1.091 cm³; when melted it becomes 1 cm³ of water. Ice contracting on melting matters for the freezing of lakes and pipes.

11.16 EFFECT OF PRESSURE ON THE MELTING POINT

Two effects appear in this section, and both are asked in the exercises.

Pressure: substances that contract on melting, like ice, have a melting point that decreases as pressure increases. Ice’s melting point drops by 0.0072°C per atmosphere of added pressure. Substances that expand on melting, like wax and lead, have a melting point that increases.

Impurity: melting point decreases with impurities. Salt mixed with ice in proper proportion lowers the melting point from 0°C to about -22°C. This freezing mixture is used to set kulfis.

11.18 VAPORISATION OR BOILING

Vaporisation is the change from liquid to vapour at a constant temperature on absorbing heat. The boiling point is that constant temperature. Condensation (or liquefaction) is the reverse change with rejection of heat, and its constant temperature is the condensation point.

Two notes. For a pure substance, the boiling point and condensation point are identical. For the same mass, the heat absorbed during vaporisation equals the heat rejected during condensation.

11.19 HEATING CURVE FOR WATER

The experiment: water at about 20°C in a flask with a suspended thermometer. Heat slowly until boiling starts. Bubbles form throughout the water, which shows boiling occurs through the whole volume.

Plot temperature against time. Part AB rises as the liquid warms. Part BC stays horizontal at 100°C while boiling continues even though heat is still absorbed. The flat line at B is the boiling point of water.

11.20 CHANGE IN VOLUME ON BOILING

All liquids expand on boiling. One cubic centimetre of water at 100°C becomes 1760 cm³ of steam at 100°C. This huge expansion is why steam can do work and why pressure builds in closed vessels.

11.21 EFFECT OF PRESSURE ON THE BOILING POINT

The boiling point of a liquid rises with increased pressure and falls with decreased pressure. Pure water boils at 100°C at one atmosphere, which is 760 mm of Hg.

Pressure cooker: trapped steam raises the pressure to about 1.75 atmospheres, so water boils at about 120°C to 125°C. Cooking becomes faster because the food gets sufficient heat energy. The cooked food also stays warm longer.

High altitude: atmospheric pressure is low, so water boils below 100°C. It gives less heat to the food, so cooking is difficult and slow.

11.22 EFFECT OF IMPURITIES ON THE BOILING POINT

Impurities raise the boiling point of a liquid. Water with common salt added boils above 100°C. This is why salt is added while cooking pulses: the salted water gives the contents sufficient heat energy, so cooking is easier and faster.

The contrast is worth memorising: pressure lowers the melting point of ice but raises the boiling point of water, and impurities do the same in each case.

11.23 LATENT HEAT AND SPECIFIC LATENT HEAT

Latent heat is the heat absorbed or liberated during a change of phase at constant temperature. It is not shown by any rise or fall in temperature. Specific latent heat is latent heat for unit mass, denoted L:

L = \frac{Q}{m}

So the heat for a given mass is Q = mL. The SI unit is J kg⁻¹. Other units are cal g⁻¹ and kcal kg⁻¹, related as 1 kcal kg⁻¹ = 1 cal g⁻¹ = 4.2 J g⁻¹ = 4.2 × 10³ J kg⁻¹.

Worked pattern: melting 0.5 kg of paraffin wax uses L = 146 \times 10^3\ \text{J kg}^{-1}, so Q = 0.5 \times 146000 = 73000\ \text{J}. The constant-temperature condition is what makes latent heat hidden from a thermometer.

11.24 SPECIFIC LATENT HEAT OF FUSION OF ICE

Two mirror definitions. Specific latent heat of melting: heat to melt unit mass of ice at 0°C into water at 0°C. Specific latent heat of fusion: heat released when unit mass of water at 0°C freezes into ice at 0°C. For a pure substance, the two values are equal.

The ice value: 336000\ \text{J kg}^{-1} = 80\ \text{cal g}^{-1}. One kilogram of ice at 0°C absorbs 336000 J to become water at 0°C. One gram of water at 0°C holds 336 J more heat than one gram of ice at 0°C.

Worked pattern: melting 150 g of ice at 0°C needs 150 \times 336 = 50400\ \text{J}.

Edge case: hot metal may not melt all the ice. Take 300 g of copper pieces at 400°C with c = 0.42\ \text{J g}^{-1}\text{K}^{-1}, mixed with 250 g of crushed ice at 0°C with L = 336\ \text{J g}^{-1}. Heat from copper cooling to 0°C is 300 \times 0.42 \times 400 = 50400\ \text{J}. Melting all the ice needs 250 \times 336 = 84000\ \text{J}.

Since 50400 is less than 84000, only 50400 / 336 = 150\ \text{g} of ice melts. The mixture stays at 0°C with ice remaining.

11.25 EXPLANATION OF LATENT HEAT OF MELTING ON THE BASIS OF KINETIC MODEL

In a solid, molecules vibrate about their mean positions. Total molecular energy is the sum of kinetic energy, which depends on temperature, and potential energy, which depends on molecular attraction and separation. During melting at constant temperature, the average kinetic energy does not change, but the average separation between molecules increases.

The supplied heat goes into that potential energy increase. That is the latent heat of melting. When 1 g of ice at 0°C changes to water at 0°C, about 336 J is used in increasing the intermolecular separation.

11.26 NATURAL CONSEQUENCES OF HIGH SPECIFIC LATENT HEAT OF FUSION OF ICE

Five consequences follow from ice’s high specific latent heat of fusion.

1. Snow on mountains melts slowly. Ice needs 336000 J kg⁻¹, so sudden floods are avoided.
2. Water in lakes and ponds freezes gradually. The ice layer formed is a poor conductor that slows further heat loss.
3. Ice at 0°C cools drinks faster than ice-cold water at 0°C. Each gram absorbs 336 J to melt.
4. When a frozen lake starts melting, the surroundings turn cold because melting draws a large amount of heat from the atmosphere.
5. It feels colder after a hailstorm because the melting ice absorbs heat from the surroundings.

What to practise from this chapter

For the calorimetry class 10 ICSE chapter, the practice blocks in Selina are:

• EXERCISE-11(A): 29 questions covering heat, temperature, heat capacity and specific heat capacity
• MULTIPLE CHOICE TYPE: 16 questions after Exercise 11(A)
• NUMERICALS: 13 numericals closing Part A
• Short-answer questions in Part B: 29 questions on phase change and latent heat
• MULTIPLE CHOICE TYPE: 14 questions near the chapter end
• NUMERICALS: 12 numericals closing Part B

Exam pattern mapping. Definition questions draw on sections 11.1 to 11.8. Application reasons draw on 11.10, 11.11 and 11.26. Graph reading draws on 11.14 and 11.19. Numericals draw on 11.3, 11.6, the 11.9 Principle, and 11.23-11.24.

The full Class 10 Physics book list and the complete Class 10 book list will help you plan the rest of your study. Verify the current syllabus on the official CISCE website.

Frequently asked questions

What is calorimetry class 10 ICSE?

Calorimetry is the measurement of the quantity of heat. In the Class 10 ICSE syllabus, it covers how heat flows between bodies, the units of heat, and the equations that connect heat, mass, temperature rise and specific heat capacity.

Why does temperature stay constant while ice melts in this ICSE chapter?

During a phase change, the supplied heat increases the potential energy of the molecules, not their kinetic energy. Since temperature depends on average kinetic energy, it stays constant until the phase change is complete.

How do you find the final temperature when ice is added to warm water in Class 10 ICSE?

Set heat lost by the water equal to the heat needed to melt the ice plus the heat to warm the melted ice water. First check whether the water can supply enough heat to melt all the ice. If not, the mixture stays at 0°C with ice remaining.

What is the difference between heat capacity and specific heat capacity in this ICSE chapter?

Heat capacity is the heat to raise the whole body by 1°C, in J K⁻¹. Specific heat capacity is the heat to raise unit mass by 1°C, in J kg⁻¹ K⁻¹. Heat capacity depends on mass and substance; specific heat capacity depends only on the substance.

Why does ice cool a drink faster than ice cold water in the Class 10 ICSE syllabus?

Each gram of ice at 0°C absorbs 336 J to melt into water at 0°C. Ice-cold water absorbs no latent heat. The extra 336 J per gram makes ice the stronger coolant.

The syllabus this chapter follows is set by the CISCE board.

Reference: ICSE Class 10 Physics textbook material, chapter 11: Calorimetry — CISCE-aligned.

More from this book

• Chapter 1: Force
• Chapter 2: Work Energy and Power
• Chapter 3: Machines
• Chapter 4: Refraction of Light At Plane Surfaces
• Chapter 5: Refraction Through a Lens
• Chapter 6: Spectrum
• Chapter 7: Sound
• Chapter 8: Current Electricity
• Chapter 9: Household Circuits
• Chapter 10: Electro Magnetism
• Chapter 12: Radioactivity