Chapter 1 — ଆମ ଚତୁଃପାର୍ଶ୍ୱରେ ଥିବା ପଦାର୍ଥ

Anything that has mass and occupies space (has volume) is called matter. A grain of sand, a drop of water, the air in a football, a mountain and a star are all matter. Heat, light, sound, electricity, love and thought are not matter — they have no mass and occupy no space, though they are real. Since ancient times Indian philosophers classified matter into five elements, the panch tatva — air, earth, fire, sky and water — and Greek thinkers had a similar idea. Modern science classifies matter by its physical state and by its chemical composition; this chapter deals with the physical nature.

The central idea is that matter is made up of particles. This was not obvious. For a long time one school held that matter was continuous, like a block of butter that can be divided without limit. The Indian philosopher Maharishi Kanad, more than 2,500 years ago, proposed that dividing matter (padarth) repeatedly would end in an indivisible particle he called parmanu. Modern experiments confirm that matter is particulate and that the particles are extremely small.

A simple activity shows both facts. Take 100 mL of water in a beaker and dissolve a spoon of salt or sugar; the level of the water does not rise noticeably. The dissolved particles have entered the spaces between the water particles. Now take a few crystals of potassium permanganate and dissolve them in 100 mL of water; take 10 mL of that deep purple solution and dilute it to 100 mL, and repeat five or six times. Even after dilution by a factor of a million the water is still faintly coloured. A few crystals therefore contain millions upon millions of particles, each far too small to be seen. The same is shown by a drop of Dettol or a few crystals of copper sulphate spreading colour through a whole glass of water.

How small are the particles? A single crystal of potassium permanganate of a milligram contains about 10¹⁸ formula units. The particles of matter are so small that even the most powerful optical microscope cannot show them; special electron microscopes can barely resolve the largest of them. Everything that follows in this chapter — the states of matter, their properties and their changes — is explained by asking three questions: how far apart are the particles, how fast are they moving, and how strongly do they attract one another.

The particles of matter have three characteristics, and each can be demonstrated by a simple activity.

1. The particles of matter have spaces between them. When sugar, salt, or potassium permanganate dissolves in water, its particles slip into the spaces between the water particles, which is why the volume does not rise and the substance seems to disappear. Gases have the largest spaces, which is why they can be compressed so much; solids have the smallest.

2. The particles of matter are continuously moving. Particles are never at rest; they possess kinetic energy. Put an unlit incense stick in a corner of a room and one has to go close to smell it; light it and the smell reaches the far corner within minutes — the hot particles of scented vapour move faster and spread through the air. Place a crystal of copper sulphate or a drop of ink at the bottom of a glass of still water and, without stirring, the colour spreads upward slowly through the whole glass over hours. This spontaneous mixing of particles of two substances by their own motion is called diffusion. As temperature rises the particles gain kinetic energy and move faster, so diffusion is faster in hot water than in cold — the reason a crystal colours warm water much sooner than cold. The random zigzag dance of pollen grains or smoke particles in air, seen under a microscope, called Brownian motion, is direct evidence that the invisible particles of the medium are moving and bumping into them.

3. The particles of matter attract each other. A force of attraction holds the particles together. Its strength varies from one kind of matter to another. Try to cut a stream of water with your fingers — the stream re-joins at once, because water particles attract each other. Now try to break an iron rod, a piece of chalk and a rubber band: the iron needs enormous force, the chalk breaks easily, the rubber stretches; the attraction between particles is strongest in iron, weaker in chalk, and of a different kind in rubber. In gases the attraction is negligible, which is why a gas spreads to fill any container.

These three characteristics together give the kinetic theory of matter: the state of a substance depends on the balance between the kinetic energy of its particles, which tends to separate them, and the force of attraction, which tends to hold them together. Where attraction wins, the substance is solid; where motion wins completely, it is a gas; a liquid is the middle case. The next topic examines each state in this light.

A solid is matter that has a definite shape and a definite volume. In a solid the particles are packed very closely, usually in a regular pattern, with very little space between them. The force of attraction between the particles is strong, so they cannot move away from their positions; they only vibrate about fixed points. This explains every property of solids.

Definite shape and volume. A brick, a pen or a stone keeps its shape wherever it is placed, because its particles cannot leave their positions. Its volume is also fixed. Rigidity. A solid resists a change of shape; the property of maintaining shape when an outside force is applied is called rigidity, and solids are rigid. Incompressibility. Since there is very little space between the particles, a solid can hardly be compressed; hitting a steel block with a hammer does not reduce its volume. High density. Closely packed particles mean a large mass in a small volume, so solids generally have higher densities than liquids and gases. Negligible fluidity. A solid does not flow. Very slow diffusion. Because the particles are held in place, diffusion in solids is extremely slow — but it does happen: two polished blocks of gold and lead pressed together for years show atoms of each in the other, and a brass chain tied tightly around a wooden post and left for years leaves a mark of metal in the wood.

Several everyday cases seem to contradict the definition and need to be understood. A rubber band changes shape when stretched, but it regains its shape when released, and it breaks if stretched too far; it is a solid. Sugar and salt take the shape of the vessel they are put in, but look at a single crystal: it keeps its own shape; it is the heap that flows, not the particles. A sponge can be compressed, but only because it has minute holes filled with air that is squeezed out; the solid material of the sponge itself is not compressed. A kinetic-energy comparison: the particles of a solid have the least kinetic energy of the three states at a given temperature.

Solids are further divided into crystalline solids, in which the particles are arranged in an orderly repeating pattern and which have sharp melting points (salt, sugar, diamond, ice), and amorphous solids, in which the arrangement is irregular and which soften over a range of temperature (glass, rubber, plastic). This detail is developed in higher classes; for now the essential picture is of particles packed tight, held firm and merely vibrating.

A liquid has a definite volume but no definite shape: it takes the shape of the container in which it is kept. In a liquid the particles are still close together — nearly as close as in a solid — but the force of attraction between them is weaker, and their kinetic energy is greater, so they can slide over one another and move about within the liquid. This freedom of movement is the key to all the properties of liquids.

Fluidity. Because the particles can move past each other, a liquid flows and is called a fluid. Water poured from a jug takes the shape of the glass; milk poured on a table spreads. The ease of flow differs from one liquid to another: water flows easily, honey and oil flow slowly because their particles attract each other more strongly (this resistance to flow is called viscosity). Definite volume. Although the shape changes, the volume does not: 250 mL of water is 250 mL in a bottle, a bowl or a balloon. Almost incompressible. The particles are close together, so a liquid cannot be compressed much; this is why brakes and hydraulic jacks use oil. Density. Liquids are generally slightly less dense than the corresponding solids (ice is the exception) and far denser than gases. Diffusion. Diffusion in liquids is faster than in solids because the particles move freely, but slower than in gases; a drop of ink in still water takes minutes to spread. Gases, liquids and solids can all diffuse into liquids: oxygen and carbon dioxide from the air dissolve in water, and this dissolved oxygen is what fish and other aquatic animals breathe.

Surface and level. A liquid at rest has a horizontal free surface. Because the particles at the surface are attracted inward by the particles below, the surface behaves like a stretched skin — surface tension — which lets a needle float on water and insects walk on a pond.

Comparison with solids. Solids and liquids are together called the condensed states because their particles are in close contact. The difference is order and freedom: in a solid the particles have fixed positions and only vibrate; in a liquid they vibrate, rotate and translate, changing neighbours constantly. The kinetic energy of the particles of a liquid at a given temperature is greater than that of the solid and less than that of the gas. This is why melting a solid needs heat and why liquids are the state in which chemical reactions and life happen most readily: particles are close enough to meet and free enough to move.

A gas has neither a definite shape nor a definite volume: it takes the shape and the whole volume of any container it is put in. In a gas the particles are very far apart compared with their own size — the spaces between them are hundreds of times their diameter — the force of attraction between them is negligible, and they move at very high speed in all directions in straight lines, colliding with each other and with the walls of the container. A gas is almost all empty space.

High compressibility. Because of the large spaces, the particles of a gas can be pushed much closer together. This is used every day: LPG (liquefied petroleum gas) in the kitchen cylinder, CNG (compressed natural gas) in vehicle tanks, oxygen in hospital cylinders, and the air in a tyre or a football are all gases compressed into a small volume. A large volume of gas can be stored and transported in a small cylinder.

No definite shape or volume; complete fluidity. A gas expands to fill any vessel, however large, because there is no attraction to hold the particles together. Open a gas cylinder in a room and the gas is soon everywhere.

Very low density. A litre of air at room temperature has a mass of only about 1.2 g, compared with 1,000 g for a litre of water. This is why gases rise in liquids as bubbles and why a hot-air balloon floats.

Very fast diffusion. Gas particles move at hundreds of metres per second, so gases mix rapidly. The smell of perfume opened in one corner of a room, the aroma of food from the kitchen, or the smell of a leaking LPG cylinder — to which a strong-smelling substance, ethyl mercaptan, is deliberately added — reaches every part of the house within minutes. A lighter gas diffuses faster than a heavier one.

Pressure. The particles of a gas hit the walls of their container continuously and in enormous numbers; the force of these collisions on each unit area of the wall is the pressure of the gas. Pump more air into a tyre and there are more particles to collide, so the pressure rises; heat a gas and its particles move faster and hit harder, so the pressure rises — which is why a sealed can must never be thrown into a fire. Atmospheric pressure, about 1 atmosphere or 101,325 pascals at sea level, is the weight of the air above us, and it supports a column of mercury 76 cm high in a barometer.

In the gaseous state the particles have the maximum kinetic energy of the three states. The whole behaviour of a gas — filling its container, being squeezed, mixing fast, pressing on its walls — follows from particles that are free, fast and far apart.

The three states are best compared side by side, because the examination often asks for the differences and the reasons behind them.

Two words need exact definition. Rigidity is the property of matter by which it keeps its shape when a force is applied; solids are rigid. Fluidity is the ability to flow; liquids and gases have it and are together called fluids. The question sometimes appears as: is a sponge or a rubber band a solid, and is a gas a fluid? Both a sponge and a rubber band are solids (their particles have fixed positions; the sponge's air holes and the rubber's elasticity explain the apparent exceptions), and gases are fluids because they flow.

Diffusion deserves a closer look since it appears in all three states. Diffusion is the intermixing of particles of two different kinds of matter on their own, due to the motion of their particles. Its rate depends on the state (gas fastest, solid slowest) and on temperature (faster when hot). Gases diffuse into gases (perfume in air), liquids into liquids (ink in water), solids into liquids (sugar into water, copper sulphate into water), gases into liquids (oxygen and carbon dioxide into water, on which aquatic life and carbonated drinks depend), and even solids into solids very slowly. Diffusion is why the smell of a hot sizzling dish reaches us from across a restaurant while a cold dish must be brought to the nose; the hot food's vapours have more kinetic energy.

Finally, a note on the ordering of states by density: a substance is normally densest as a solid, less dense as a liquid and least dense as a gas, because the particles move apart as the kinetic energy increases. Water is the celebrated exception — ice floats — and this single anomaly keeps lakes from freezing solid and lets fish survive winter under a floating lid of ice.

A substance can be changed from one state to another by changing its temperature or the pressure on it. We begin with heating a solid.

Take about 150 g of crushed ice in a beaker, insert a thermometer so that the bulb is in the ice, and heat gently on a low flame, stirring. The temperature rises until the ice begins to melt, and then, though heat is still being supplied, the thermometer stays at 0 degrees Celsius until the last piece of ice has melted. Only after that does the temperature of the water begin to rise again. The temperature at which a solid melts to become a liquid at atmospheric pressure is its melting point; for ice it is 0 degrees Celsius, which on the Kelvin scale is 273.15 K (often written 273 K). The melting point is a measure of the strength of the force of attraction between the particles of a solid: iron melts at 1,538 degrees Celsius, wax at about 60, and ice at 0.

What is happening at the particle level? As the solid is heated, its particles gain kinetic energy and vibrate more violently. At the melting point the vibration becomes strong enough to overcome the forces of attraction; the particles leave their fixed positions and become free to move — the solid has become a liquid. The process of a solid changing to a liquid on heating is called melting or fusion.

Why does the temperature not rise during melting although heat is being supplied? Because the heat is being used to break the attractions between the particles, not to increase their speed. This hidden heat is called the latent heat of fusion: the amount of heat energy required to change 1 kg of a solid into liquid at its melting point at atmospheric pressure, without any change of temperature. For ice it is 3.34 × 10⁵ joules per kilogram (about 80 calories per gram). This means that water at 0 degrees Celsius has 3.34 × 10⁵ J more energy per kilogram than ice at the same temperature — and this is why ice at 0 degrees Celsius is more effective at cooling a drink than water at 0 degrees Celsius: the ice absorbs its latent heat from the drink as it melts.

The reverse process, a liquid turning into a solid on cooling, is freezing or solidification, and it happens at the same temperature, the freezing point; the latent heat is released to the surroundings during freezing. Farmers in cold countries spray water on fruit trees before a frost for exactly this reason: as the water freezes it gives out latent heat and protects the buds.

Continue heating the water obtained by melting the ice. Its temperature rises steadily as the particles gain kinetic energy and move faster. At 100 degrees Celsius bubbles of vapour form throughout the liquid, rise and burst at the surface — the water is boiling — and again the thermometer stops rising, staying at 100 degrees Celsius until all the water has turned into steam. The temperature at which a liquid starts boiling at atmospheric pressure is its boiling point. For water it is 100 degrees Celsius, or 373 K. Boiling is a bulk phenomenon: particles from the whole of the liquid, not only from the surface, gain enough energy to escape into the gaseous state.

At the particle level, at the boiling point the kinetic energy of the particles becomes enough to overcome the forces of attraction completely; the particles separate from each other entirely and fly off as a gas. As with melting, the temperature stays constant during boiling because the supplied heat is used to break the attractions rather than to speed the particles up. This heat is the latent heat of vaporisation: the heat energy required to change 1 kg of a liquid into gas at its boiling point at atmospheric pressure. For water it is 22.6 × 10⁵ J/kg (about 540 calories per gram) — nearly seven times the latent heat of fusion, because vaporisation must separate the particles completely, while melting only loosens them.

This large latent heat has a serious practical consequence: steam at 100 degrees Celsius causes far more severe burns than boiling water at 100 degrees Celsius. When steam touches the skin it condenses, releasing 22.6 × 10⁵ J per kg in addition to the heat that the resulting hot water gives out as it cools. This is a favourite examination question.

The reverse of boiling is condensation — a gas changing into a liquid on cooling — and it releases the latent heat of vaporisation. Dew on grass in the early morning, drops on the outside of a glass of iced water, and clouds are condensation of water vapour. In a laboratory a condenser cools vapour back to liquid in distillation.

The boiling point, like the melting point, depends on the strength of attraction between particles and is a characteristic of the substance: ethanol boils at 78 degrees Celsius, water at 100, mercury at 357. It also depends on pressure, which is treated in a later topic. A summary of the terms: the change of a solid to liquid is fusion, of liquid to gas is vaporisation, of gas to liquid is condensation, of liquid to solid is freezing; the heat exchanged in each without a change of temperature is latent heat.

Some substances do not pass through the liquid state at all when heated: they change directly from solid to gas. This is sublimation, and the reverse change, from gas directly to solid, is called deposition (or also sublimation in some books). The substances that behave this way have particles in the solid that are held only weakly, so that, at atmospheric pressure, the energy that would melt them is enough to set them completely free.

The classic demonstration uses ammonium chloride (nausadar). Place a little in a china dish, cover it with an inverted funnel whose stem is plugged with cotton, and heat gently. The white powder disappears from the dish without melting, and white crystals appear on the cool inner wall of the funnel: the solid sublimed to vapour, rose, and deposited on the cold surface. Camphor (kapur) and naphthalene balls sublime slowly even at room temperature, which is why naphthalene balls kept among woollens shrink and vanish over months, leaving no liquid behind. Iodine crystals heated in a test tube give a beautiful violet vapour that re-forms crystals on the cool upper wall. Solid carbon dioxide, called dry ice, sublimes at minus 78 degrees Celsius at atmospheric pressure and is used to keep ice-cream frozen and to make stage fog; it is called dry because it never wets — it leaves no liquid. Sublimation is used to purify substances such as camphor and to separate a subliming component from a non-subliming one, such as ammonium chloride from common salt.

The changes of state can now be drawn together into one picture. On heating, and with increasing kinetic energy of the particles: solid → liquid (fusion) → gas (vaporisation); or solid → gas directly (sublimation). On cooling, and with decreasing kinetic energy: gas → liquid (condensation) → solid (freezing or solidification); or gas → solid directly (deposition). At every change the substance remains chemically the same — ice, water and steam are all H₂O — so change of state is a physical change, reversible by reversing the temperature or pressure. This is different from a chemical change such as the burning of paper, which cannot be undone by cooling.

The temperatures and latent heats that mark these changes are fixed for each pure substance at a given pressure, and they serve as fingerprints: a liquid that boils at exactly 100 degrees Celsius at sea level is pure water, and a sample that boils over a range is a mixture — an idea that the next chapter uses to test purity.

Temperature is not the only way to change the state of matter. Since the difference between the states is the distance between particles, pressure — which pushes particles closer — can also change the state, and this is most easily seen with gases.

Take a gas in a cylinder fitted with a piston. When the piston is pushed in, the volume of the gas decreases and its particles are brought closer together. If enough pressure is applied, and especially if the gas is also cooled, the particles come close enough for the forces of attraction to hold them, and the gas becomes a liquid. This is liquefaction. Both cooling and compression favour it: cooling reduces the kinetic energy that keeps the particles apart, and pressure reduces the space between them.

This principle is used in every kitchen. LPG — mainly butane and propane — is a gas at room temperature and atmospheric pressure, but inside the cylinder it is kept under a pressure of a few atmospheres and is a liquid; if you shake a cylinder you can hear it slosh. When the valve is opened the pressure falls, the liquid boils and comes out as gas to the burner. Similarly CNG in vehicles, oxygen and nitrogen in hospital and industrial cylinders, and the refrigerant gas in a refrigerator or air conditioner are compressed. The refrigerator, in fact, works by a cycle of compressing a gas (which liquefies and gives out heat at the back) and then letting the liquid expand and evaporate inside (which absorbs heat from the food).

Dry ice illustrates the effect of pressure on sublimation. At atmospheric pressure solid carbon dioxide passes directly to gas at minus 78 degrees Celsius; liquid carbon dioxide can exist only above a pressure of about 5 atmospheres. So carbon dioxide is stored as a liquid in cylinders under high pressure, and when it is released into the open air it turns instantly into solid dry ice and gas. The state of a substance thus depends on both temperature and pressure, and a diagram of pressure against temperature, called a phase diagram, shows the region of each state; the student meets it in later classes.

Pressure also affects melting and boiling points. Increasing the pressure raises the boiling point of a liquid: in a pressure cooker the steam cannot escape, the pressure rises to about 2 atmospheres, water boils at about 120 degrees Celsius, and food cooks faster. Decreasing the pressure lowers the boiling point: on a hill station the air pressure is lower, water boils below 100 degrees Celsius, and rice and dal take much longer to cook. For ice, an increase of pressure slightly lowers the melting point, which is why a wire loaded with weights can pass slowly through a block of ice, the ice re-freezing behind it (regelation), and why ice skates glide on a thin film of melted water.

The changes of state are described by temperatures, so the scales used to measure temperature must be understood. Temperature is a measure of the average kinetic energy of the particles of a substance — the hotter a body, the faster its particles move. It is measured with a thermometer.

The Celsius scale (formerly centigrade) sets the melting point of ice at 0 degrees Celsius and the boiling point of water at 100 degrees Celsius at one atmosphere, and divides the interval into 100 equal degrees. It is the scale of daily life, of weather reports and of the clinical thermometer (normal body temperature about 37 degrees Celsius).

The Kelvin scale is the SI scale of temperature, named after Lord Kelvin. Its unit is the kelvin, symbol K, written without a degree sign. Its zero, called absolute zero, is the lowest temperature possible — the temperature at which the particles of matter would have their minimum possible energy — and equals minus 273.15 degrees Celsius. The size of one kelvin is the same as one degree Celsius, so the two scales differ only in where they start. To convert: temperature in kelvin = temperature in degrees Celsius + 273.15, and for school work we use 273. Thus 0 degrees Celsius = 273 K, 25 degrees Celsius = 298 K, 100 degrees Celsius = 373 K, and 300 K = 27 degrees Celsius. There are no negative temperatures on the Kelvin scale, which is why it is the scale used in scientific equations.

The Fahrenheit scale, used in some countries and in older clinical thermometers, sets the ice point at 32 and the steam point at 212 degrees Fahrenheit; body temperature is 98.6 degrees Fahrenheit. The conversion is F = (9/5) C + 32, but the Odisha examination deals almost entirely with Celsius and Kelvin.

A few points of care. A change in temperature is the same number in Celsius and Kelvin: if water warms from 20 to 30 degrees Celsius, it warms by 10 degrees Celsius, which is also a rise of 10 K. The Kelvin scale is called an absolute scale because it starts from absolute zero. In the chapter's data, the melting point of ice is 273 K and the boiling point of water 373 K; it is good practice to give both values in the examination when a question asks for the temperature of a change of state.

A convenient table of the important temperatures of this chapter: melting point of ice 0 degrees Celsius = 273 K; boiling point of water 100 degrees Celsius = 373 K; sublimation of dry ice minus 78 degrees Celsius = 195 K; normal room temperature 25 degrees Celsius = 298 K; absolute zero minus 273 degrees Celsius = 0 K.

Water boils at 100 degrees Celsius, yet wet clothes dry, puddles disappear and water in an open dish slowly vanishes at ordinary temperatures. The change of a liquid into vapour at any temperature below its boiling point is called evaporation.

How does it happen? The particles of a liquid at a given temperature do not all have the same kinetic energy; there is a range, and a small fraction at the surface have, at any moment, enough energy to break free of the attraction of their neighbours and escape into the air. Evaporation is therefore a surface phenomenon, unlike boiling, which happens throughout the liquid. It goes on at all temperatures, silently and without bubbles, and it is faster when more particles have the energy to escape and when the escaped particles are carried away.

Four factors control the rate of evaporation.

1. Surface area. Since evaporation occurs at the surface, a larger surface means a faster rate. This is why clothes are spread out to dry rather than left in a heap, why tea is poured into a saucer to cool it quickly, and why the same water in a plate dries faster than in a glass.

2. Temperature. At a higher temperature more particles have enough kinetic energy to escape, so the rate increases. Clothes dry faster on a hot day, and a wet floor dries quickly in May but stays damp in a January morning.

3. Humidity. Humidity is the amount of water vapour already present in the air. Air can hold only a certain amount of vapour at a given temperature; when it is nearly saturated, evaporation is slow. This is why clothes take so long to dry in the humid monsoon weeks of July in coastal Odisha, and why the same weather feels sticky — our sweat cannot evaporate.

4. Wind speed. Moving air carries away the vapour particles from near the surface, making room for more to escape. Clothes dry faster on a windy day; blowing on hot tea cools it faster; a fan makes us feel cool.

A summary: rate of evaporation increases with surface area, temperature and wind speed, and decreases with humidity. Two more points are often asked. Evaporation is not the same as boiling: boiling occurs at one fixed temperature, throughout the liquid, rapidly and with bubbles; evaporation occurs at all temperatures, only at the surface, slowly and silently. And the vapour of a liquid at room temperature is called a vapour rather than a gas because the substance is a liquid at that temperature; the word gas is used for a substance that is gaseous at room temperature.

Evaporation has a consequence that we feel on our skin every day: it causes cooling. The reason lies in which particles escape. Only the most energetic particles at the surface have enough kinetic energy to break away; when they leave, the average kinetic energy of the particles remaining in the liquid falls, and since temperature is a measure of average kinetic energy, the temperature of the liquid falls. The liquid then absorbs heat from its surroundings — the vessel, the skin, the air — to make up the loss, and the surroundings become cool. In other words the latent heat of vaporisation is taken from whatever the liquid is in contact with.

This single principle explains a long list of everyday observations that the examination loves to ask.

Sweating. On a hot day or after exercise, the body sweats; the sweat evaporates from the skin, taking the latent heat from the body, and we feel cool. On a humid day sweat cannot evaporate, and we feel hot and sticky; a fan helps because moving air speeds evaporation.

Water in an earthen pot (matka or surahi). The clay has tiny pores through which water seeps to the outer surface and evaporates continuously, drawing heat from the water inside, which stays several degrees cooler than room temperature. A glazed or metal pot has no pores and does not cool its water.

Cotton clothes in summer. Cotton absorbs sweat and exposes it to the air over a large surface, so it evaporates quickly and cools the body; synthetic fabrics do not absorb sweat. White or light colours also reflect heat.

Acetone, petrol or spirit on the palm. These liquids evaporate very quickly because their particles are weakly attracted; they take the latent heat from the palm, which feels distinctly cold. This is why a doctor's spirit swab feels cold before an injection.

Sprinkling water on a roof or open ground in the evening of a hot day: the water evaporates, taking the latent heat from the hot surface, and the room or courtyard cools.

Desert coolers and the wet-cloth trick. A desert cooler draws hot dry air through pads kept wet with water; evaporation from the pads cools the air. It works well in the dry heat of western India but poorly in the humid coastal air of Odisha, where an air conditioner is needed instead. A bottle of water wrapped in a wet cloth and kept in a breeze stays cool for the same reason. Dogs pant and elephants flap their ears to increase evaporation. The refrigerator, at bottom, is an evaporation machine: the liquid refrigerant evaporates inside and takes heat from the food.

The student should be able to reverse the reasoning too: condensation causes heating, which is why a room feels warmer just before rain and why steam burns so severely. Evaporation and condensation are the two directions of one exchange of latent heat.

Solid, liquid and gas are the three states we meet every day, but scientists recognise two more, which the chapter introduces briefly so that the student knows the picture is larger.

Plasma is the fourth state of matter. If a gas is heated to a very high temperature, or subjected to a strong electric field, the collisions between its particles become violent enough to knock electrons out of the atoms. The gas then consists of positively charged ions and free electrons — a mixture that is electrically neutral as a whole but conducts electricity and glows. Such an ionised gas is called plasma. The glow inside a fluorescent tube and a neon sign is plasma: the electric current ionises the gas (mercury vapour or neon) inside the tube and the plasma glows with a colour that depends on the gas — orange-red for neon, blue-green for argon with mercury. Lightning is plasma. The Sun and all the stars are enormous balls of plasma, kept hot by nuclear fusion; in fact more than 99 percent of the visible matter in the universe is in the plasma state, so from the universe's point of view plasma is the ordinary state and our solids and liquids are the rare exceptions. Plasma is also used in plasma televisions, in cutting and welding torches, and in the experimental fusion reactors that hope to copy the Sun.

Bose-Einstein condensate (BEC) is the fifth state, and it carries an Indian name. In 1924 the Indian physicist Satyendra Nath Bose sent Albert Einstein a paper on the statistics of light particles; Einstein extended the idea to atoms and predicted that if a gas of certain atoms were cooled to an extremely low temperature — within a few billionths of a degree of absolute zero — at a very low density, the atoms would lose their individual identity and collapse into a single quantum state, behaving like one giant atom. This state, named the Bose-Einstein condensate in honour of both scientists, was actually produced in 1995 by Eric Cornell, Carl Wieman and Wolfgang Ketterle, who cooled rubidium atoms with lasers and magnetic traps; they received the Nobel Prize in Physics in 2001. The BEC exists only in the laboratory and is used to study the strange laws of the quantum world.

The five states therefore run from the coldest to the hottest: Bose-Einstein condensate (near absolute zero), solid, liquid, gas, and plasma (thousands of degrees or more). The three that the chapter studies in detail are the states at the temperatures and pressures of ordinary life. What unites all five is the principle with which the chapter began: matter is made of particles, and the state depends on how much energy those particles have and how strongly they hold one another.