Matter in Our Surroundings

1. What is Matter?

Ancient Indian philosophers classified everything in the universe into five elements called Panch Tatva: air, earth, fire, sky, and water. Modern science defines matter precisely as anything that has mass and occupies space.

Examples of matter: water, rice, soil, iron, air, your body. Counter-examples (not matter): light, heat, sound, gravity, emotions.

The particles of matter are so tiny that they are invisible to the naked eye. One drop of ink dissolved in a glass of water can colour the whole glass — evidence that ink particles are incredibly small and can spread everywhere.

2. Physical Nature of Matter — Particles, Spaces, Motion

NCERT establishes three key ideas about the particle nature of matter through simple activities.

2.1 Particles have inter-particle spaces

Mix 50 mL of water and 50 mL of alcohol. The total volume is less than 100 mL. Why? The particles of alcohol fit into the spaces between water particles — proof that there are spaces between particles. Similarly, dissolving salt or sugar in water without increasing the water level confirms that solute particles occupy the empty inter-particle spaces.

2.2 Particles are in constant motion

Place a crystal of KMnO4 at the bottom of a beaker of still water without stirring. After some time, the whole water turns purple. This spreading of particles on their own is called diffusion. It happens because particles are in constant, random motion.

The smell of perfume or food spreads across a room because gas-phase particles move rapidly and fill the available space — another example of diffusion in gases.

2.3 Particles attract each other

There is a force of attraction between the particles of matter called intermolecular (interparticle) force of attraction. This force holds the particles together. Its strength determines the state of matter:

• Solids — very strong attraction; particles tightly packed.
• Liquids — moderate attraction; particles can slide past one another.
• Gases — negligible attraction; particles move freely and far apart.

3. Characteristics of Particles of Matter

The five fundamental characteristics summarised:

• Very small size — a single particle cannot be seen with the naked eye or even a standard optical microscope.
• Continuous motion — particles never stop moving. Higher the temperature, higher the kinetic energy, faster the motion. At absolute zero (-273 degrees C) motion theoretically ceases.
• Intermolecular spaces — gaps exist between particles; smallest in solids, moderate in liquids, very large in gases.
• Intermolecular forces — attractive forces keep particles together; strongest in solids, weakest in gases.
• Intermingling — particles of one substance can mix with particles of another, leading to diffusion.

4. States of Matter

4.1 Solid State

In solids, particles are very closely packed in an orderly arrangement. The intermolecular forces are very strong, so particles can only vibrate about their fixed positions — they cannot move freely. Because there is almost no free space, solids cannot be compressed. Examples: iron, wood, rock, ice, chalk.

Note on exceptions: Rubber is elastic (can be stretched) but is still a solid because it returns to its original shape. Sand and salt are solids — each grain is a solid particle, though the bulk material can appear to "flow" loosely when poured.

4.2 Liquid State

In liquids, particles are close but not as tightly packed as in solids. They have enough kinetic energy to slide past one another, giving liquids their fluidity. Inter-particle spaces are slightly larger than in solids, so liquids are very slightly compressible. Examples: water, milk, oil, mercury. Liquids exert pressure in all directions and have a free surface — the top surface is flat (horizontal) due to gravity.

4.3 Gaseous State

In gases, particles are very far apart, inter-particle forces are negligible, and kinetic energy is very high. Gas particles move randomly at high speeds in all directions. Because the spaces between particles are enormous, gases are easily compressed. Examples: air, oxygen, carbon dioxide, hydrogen. Gas pressure on container walls is due to the collision of gas particles with the walls.

4.4 Comparison Table

5. Plasma and Bose-Einstein Condensate (BEC)

Beyond the three everyday states, scientists recognise two more extreme states of matter.

5.1 Plasma — The Fourth State

• Plasma is found in stars (including the Sun), lightning bolts, and the cores of flames.
• Everyday uses: neon signs, fluorescent tubes, and plasma TVs all contain plasma.
• Plasma is the fourth state of matter and is the most abundant form of visible matter in the universe.
• Unlike ordinary gases, plasma is electrically conducting and responds strongly to magnetic and electric fields.

5.2 Bose-Einstein Condensate — The Fifth State

• Predicted by Satyendra Nath Bose and Albert Einstein in 1924-25.
• First created experimentally in 1995 by Eric Cornell and Carl Wieman using rubidium atoms cooled to 170 nanokelvin.
• BEC is the fifth state of matter — exists only under extreme laboratory conditions.
• At such temperatures particles almost stop moving and lose their individual identity, behaving as one.

6. Interconversion of States of Matter

States of matter are not fixed forever — they can be changed by altering temperature or pressure.

6.1 Effect of Temperature

On heating, particles gain kinetic energy. If enough energy is supplied, the forces holding particles together are overcome and the state changes:

• Solid to Liquid: Melting (Fusion). The fixed temperature at which a solid melts is its melting point. For ice: 0 degrees C = 273 K.
• Liquid to Gas: Vaporisation (Boiling). The fixed temperature at which a liquid boils is its boiling point. For water: 100 degrees C = 373 K.
• Gas to Liquid: Condensation. On cooling, a gas loses energy; attractive forces gradually overcome kinetic energy and particles come together to form liquid.
• Liquid to Solid: Freezing (Solidification). On further cooling, particles slow down enough for strong attractive forces to lock them in a fixed arrangement.
• Solid directly to Gas: Sublimation. Some solids convert directly to vapour without passing through the liquid state. The reverse (gas to solid directly) is called deposition. Examples: dry ice (solid CO2), camphor, naphthalene (moth balls), ammonium chloride, iodine.

6.2 Effect of Pressure

Increasing pressure brings particles closer together, which can convert a gas into a liquid (or even a solid) without changing temperature. This is how LPG (liquefied petroleum gas) is stored in cylinders — the gas is compressed under high pressure until it becomes liquid, enabling large quantities to be stored compactly.

6.3 Temperature Scales and Conversion

The SI unit of temperature is the Kelvin (K). Scientists prefer Kelvin because it starts at absolute zero — there are no negative temperatures on the Kelvin scale. Absolute zero is the temperature at which particles theoretically have zero kinetic energy (minimum possible energy).

7. Latent Heat

When a substance is changing state (e.g., melting or boiling), it absorbs heat energy without any rise in temperature. This hidden heat is called latent heat (from the Latin latere meaning "to lie hidden").

7.1 Latent Heat of Fusion

When ice melts at 0 degrees C, the temperature of the ice-water mixture stays at 0 degrees C until all the ice has melted. The heat absorbed is used entirely to break the bonds between particles, not to raise temperature. This is why ice at 0 degrees C cools a drink more effectively than water at 0 degrees C — the ice absorbs latent heat from the drink as it melts, providing extra cooling.

7.2 Latent Heat of Vaporisation

When water boils at 100 degrees C, the temperature stays at 100 degrees C until all the water has converted to steam. Steam at 100 degrees C contains more energy than water at 100 degrees C by an amount equal to the latent heat of vaporisation. This is why a steam burn is much more severe than a boiling water burn at the same temperature — steam releases extra latent heat as it condenses on the skin.

8. Evaporation

Evaporation is the process by which a liquid converts into vapour at any temperature below its boiling point, from its surface only. This distinguishes it from boiling, which occurs throughout the liquid only at the boiling point.

Why does evaporation happen at all temperatures? At any temperature, some particles near the surface have enough kinetic energy to overcome the intermolecular forces and escape into the vapour phase. Only the fastest (most energetic) particles escape, which is why evaporation also causes cooling.

8.1 Factors Affecting the Rate of Evaporation

(a) Temperature: Higher temperature means particles have more kinetic energy; more particles can overcome intermolecular forces and escape. Wet clothes dry faster on a hot sunny day than on a cold day.

(b) Surface area: A larger exposed surface means more particles are at the surface and available to escape. Wet clothes spread on a clothesline dry faster than if bunched together in a heap (which has less exposed surface).

(c) Humidity (moisture content of air): If the air is already saturated with water vapour (high humidity), it cannot accept more water particles easily, so evaporation slows down. On a hot but humid day, sweat does not evaporate easily — we feel sticky and uncomfortable. On a hot and dry day, sweat evaporates quickly and we feel cooler.

(d) Wind speed: Moving air carries away the water vapour that has just evaporated, reducing the concentration of vapour above the liquid surface and allowing more evaporation to occur. Clothes dry faster on a windy day than on a still day for this reason.

9. Evaporation Causes Cooling

When a liquid evaporates, the particles that escape into the vapour phase are the most energetic (fastest-moving) ones. The remaining liquid particles have, on average, lower kinetic energy — so the temperature of the remaining liquid drops. This is the molecular explanation for why evaporation causes cooling.

Real-life applications of cooling by evaporation:

• Sweating: Sweat on the skin evaporates and absorbs latent heat from the body, keeping us cool. Dogs pant to increase evaporation from the tongue for the same reason (they have very few sweat glands).
• Earthen pot (matka): Water stored in an unglazed earthen pot stays cool because water seeps through the porous walls and evaporates from the outer surface, absorbing heat from the water inside.
• Refrigerator: A refrigerant liquid evaporates inside the refrigerator, absorbing heat from the contents; the vapour is then compressed and condensed outside, releasing the heat. The cooling is entirely due to evaporation of the refrigerant.
• Wearing cotton in summer: Cotton absorbs sweat and allows it to evaporate, keeping the body cool. Synthetic fibres do not absorb sweat as well, so sweat accumulates and does not evaporate efficiently.
• Desert cooler: Blows air over wet pads. Water evaporates from the pads, absorbing heat from the air, which then enters the room cooler. Works best in hot, dry conditions (low humidity).

Evaporation vs Boiling — Key Differences