NCERT Class 9 Science Chapter 1: Matter in Our Surroundings - A Complete Guide

Introduction to Matter in Our Surroundings

Welcome to our comprehensive guide on Chapter 1 of the NCERT Class 9 Science textbook, 'Matter in Our Surroundings'. This foundational chapter introduces you to the very substance of our universe: matter. Everything you see, touch, or feel, from the air you breathe to the book you are reading, is composed of matter. Understanding its properties and behaviours is the first step into the fascinating world of chemistry and physics. This chapter explores the physical nature of matter, its different states—solid, liquid, and gas—and how it can transition from one state to another. We will delve into concepts like particle nature, kinetic energy, intermolecular forces, melting, boiling, evaporation, and more. A solid grasp of these concepts is crucial for building a strong foundation in science. Let's begin our journey to understand the world at a molecular level!

Physical Nature of Matter

For a long time, scientists debated the nature of matter. One school of thought believed matter was continuous, like a block of wood. The other believed matter was made of discrete particles, like sand. Let's explore the evidence that proves matter is, in fact, particulate in nature.

Matter is Made Up of Particles

The idea that matter consists of tiny particles is not just a theory; it's a verifiable fact. We can demonstrate this with a simple activity.

Activity: Dissolving Salt in Water

• Procedure: Take a 100 ml beaker and fill it halfway with water. Mark the initial level of the water. Now, add one or two spoonfuls of salt (or sugar) and stir it well with a glass rod.
• Observation: You will notice that the salt disappears completely into the water. More importantly, observe the water level. The level of the water in the beaker does not rise.
• Conclusion: Where did the salt go? If matter were continuous, there would be no space to accommodate the salt particles, and the water level would have risen. The fact that the water level remains the same indicates that there are tiny, invisible spaces between the particles of water. The particles of salt break down into their own constituent particles and occupy these intermolecular spaces. This simple experiment provides strong evidence that matter is not continuous but is made up of \textremely small particles.

How Small are These Particles of Matter?

The particles that make up matter are unimaginably small—far beyond what we can see with the naked eye. To get a sense of their minuscule size, consider the following activity.

Activity: Diluting Potassium Permanganate

• Procedure: Take 2-3 small crystals of potassium permanganate (KMnO₄) and dissolve them in 100 ml of water in a beaker. You will get a deep purple solution. Now, take out approximately 10 ml of this solution and add it to another beaker containing 90 ml of clear water. This dilutes the solution.
• Observation 1: The new solution is still coloured purple, although the intensity of the colour is less than the original solution.
• Procedure Continued: Repeat this process 5 to 8 times. Each time, take 10 ml from the previous solution and add it to 90 ml of clear water.
• Observation 2: Even after multiple dilutions, the water remains coloured. This indicates that the original 2-3 crystals of potassium permanganate must have been composed of millions and millions of tiny particles. These particles kept dividing themselves and spreading throughout the water with each dilution.
• Conclusion: This activity powerfully demonstrates that the particles of matter are \textremely small. A single crystal contains millions of particles, which are capable of imparting colour to a large volume of water. The same activity can be performed with a few drops of Dettol, where the smell can be detected even after repeated dilutions.

Characteristics of Particles of Matter

The particulate nature of matter is defined by some key characteristics that govern its behaviour and properties. These characteristics explain why different substances exist as solids, liquids, or gases.

Particles of Matter Have Space Between Them

As we established from the salt-in-water experiment, there are gaps, known as intermolecular or interparticle spaces, between the particles of matter. The \textent of these spaces varies between different states of matter. In solids, these spaces are minimal. In liquids, they are larger, allowing particles to slide past one another. In gases, the spaces are vast, allowing particles to move around freely. When we make tea, coffee, or lemonade, particles of one type of matter get into the spaces between particles of the other. This property is fundamental to the concept of diffusion and solutions.

Particles of Matter are Continuously Moving

Particles of matter are never at rest; they are in constant, random motion. This motion is due to the energy they possess, known as kinetic energy. The higher the temperature, the more kinetic energy the particles have, and the faster they move. This movement is why a substance can diffuse into another.

• Example 1 (Gases): The smell of an unlit incense stick requires you to get very close to it. However, when you light the incense stick, the heat increases the kinetic energy of the aromatic particles. They move rapidly and mix with the air particles, allowing the fragrance to spread across the entire room in seconds.
• Example 2 (Liquids): Drop a crystal of copper sulphate or a drop of ink into a glass of water. Without stirring, you will observe the colour slowly and evenly spreading throughout the water over time. This happens because the particles of both the ink and water are in constant motion, causing them to intermix. This intermixing of particles of two different types of matter on their own is called diffusion. Diffusion becomes faster as the temperature increases.

Particles of Matter Attract Each Other

Particles of matter are held together by forces of attraction, known as intermolecular or interparticle forces. The strength of this force varies from one type of matter to another. This force keeps the particles bound together.

• In solids, this force is very strong, which is why solids have a fixed shape and are difficult to break. Consider trying to break a piece of chalk, an iron nail, and a rubber band. The iron nail is the hardest to break, indicating the strongest forces of attraction between its particles.
• In liquids, the force is weaker than in solids. This allows the particles to move around but keeps them within the volume of the liquid. You can easily cut through water with your hand, showing the forces are weak enough to be overcome.
• In gases, this force is negligible. The particles are far apart and move randomly in all directions. This is why gases have no fixed shape or volume and fill any container they are in.

States of Matter

Based on the variation in the characteristics of particles (interparticle space, kinetic energy, and force of attraction), matter can be classified into three main states: solid, liquid, and gas.

The Solid State

Solids are characterized by strong intermolecular forces and tightly packed particles with very little space between them. The particles vibrate about their fixed positions but do not move from place to place.

Properties of Solids:

• Definite Shape and Volume: Solids maintain their shape and volume regardless of the container they are in.
• Rigidity: Solids are rigid and resist changes in their shape. They can break under excessive force but do not flow.
• High Density: Due to closely packed particles, solids generally have high densities.
• Incompressibility: The space between particles is minimal, making solids nearly incompressible. For example, a sponge is compressible not because the solid itself is compressed, but because it has tiny pores filled with air, and it's the air that is expelled.
• No Fluidity: Solids cannot flow.

Examples: Wood, stone, iron, sugar, salt.

The Liquid State

In liquids, the intermolecular forces are weaker than in solids but strong enough to keep the particles together. The particles have more kinetic energy and can slide past one another.

Properties of Liquids:

• No Definite Shape, but Definite Volume: Liquids take the shape of the container they are in, but they have a fixed volume.
• Fluidity: Liquids can flow and are not rigid. This property is called fluidity.
• Low Compressibility: Liquids are slightly more compressible than solids but are still considered largely incompressible for most practical purposes.
• Lower Density than Solids (usually): Particles are less tightly packed than in solids, so most liquids have a lower density. A key exception is ice (solid water), which is less dense than liquid water and thus floats.
• Diffusion: Gases like oxygen and carbon dioxide can diffuse and dissolve in water. This is essential for the survival of aquatic plants and animals.

Examples: Water, milk, oil, juice.

The Gaseous State

In the gaseous state, the intermolecular forces are \textremely weak (almost negligible), and the particles are far apart. They possess high kinetic energy and move randomly and rapidly in all directions.

Properties of Gases:

• No Definite Shape or Volume: Gases take the shape and volume of their container.
• High Fluidity: Gases flow very easily.
• High Compressibility: Due to the large spaces between particles, gases can be easily compressed. This property is used in LPG (Liquefied Petroleum Gas) cylinders for cooking and CNG (Compressed Natural Gas) for vehicles.
• Low Density: Gases have very low densities because their particles are spread far apart.
• Exert Pressure: Gas particles are in constant, random motion. They collide with each other and with the walls of the container. These collisions with the walls exert a force, which, per unit area, is defined as the pressure of the gas.

Examples: Air, oxygen, hydrogen, carbon dioxide.

Can Matter Change Its State?

Yes, matter can change from one state to another. Water is the most common example: it exists as a solid (ice), a liquid (water), and a gas (water vapour). These transitions are caused by changes in temperature and pressure.

Effect of Change of Temperature

When we supply heat energy to a substance, the kinetic energy of its particles increases. This increased energy can overcome the forces of attraction between the particles, leading to a change of state.

From Solid to Liquid: Melting (Fusion)

When a solid is heated, its particles gain kinetic energy and start vibrating more vigorously. At a certain temperature, the particles gain enough energy to overcome the strong forces of attraction holding them in their fixed positions. They start to move freely, and the solid transforms into a liquid.

• Melting Point: The specific temperature at which a solid melts to become a liquid at atmospheric pressure is called its melting point. The melting point of ice is 0° Celsius (273.15 Kelvin).
• Latent Heat of Fusion: During the process of melting, you will observe that the temperature of the system does not rise, even though heat is continuously being supplied. For example, when ice is melting, the temperature remains at 0°C until all the ice has turned into water. This is because the supplied heat energy is used up in breaking the forces of attraction between the ice particles. This hidden heat, which is absorbed without causing a rise in temperature, is called the latent heat of fusion. It is defined as the amount of heat energy required to change 1 kg of a solid into a liquid at its melting point at atmospheric pressure.

From Liquid to Gas: Boiling (Vaporization)

When a liquid is heated, its particles gain even more kinetic energy. At a certain point, the particles have enough energy to break free completely from the forces of attraction of other particles. They escape from the bulk of the liquid and turn into a gas. This phenomenon is called boiling.

• Boiling Point: The temperature at which a liquid starts boiling at atmospheric pressure is known as its boiling point. The boiling point of water is 100° Celsius (373 Kelvin).
• Latent Heat of Vaporization: Similar to melting, the temperature of the water remains constant at 100°C during boiling until all the water has converted to steam. The heat energy supplied is used to overcome the intermolecular forces in the liquid state. This is called the latent heat of vaporization. It is defined as the amount of heat energy required to change 1 kg of a liquid into a gas at its boiling point at atmospheric pressure. Particles in steam at 100°C have more energy than particles in water at the same temperature because they have absorbed the \textra energy in the form of latent heat of vaporization. This is why burns from steam are often more severe than burns from boiling water.

Direct Transition: Sublimation

Some substances can change directly from a solid state to a gaseous state without passing through the liquid state. This process is called sublimation. The reverse process, changing directly from a gas to a solid, is called deposition.

• Examples: Camphor (kapur), ammonium chloride, and naphthalene balls undergo sublimation. When you heat ammonium chloride, it turns directly into a vapour, which then solidifies (deposition) on the cooler parts of the container.

Effect of Change of Pressure

Pressure can also cause a change in the state of matter. Applying pressure and reducing temperature can liquefy gases.

When we apply high pressure to a gas in a cylinder, the particles are forced closer together. The intermolecular spaces decrease significantly. If we also reduce the temperature (which lowers the kinetic energy of the particles), the intermolecular forces can become strong enough to turn the gas into a liquid. This is how LPG is stored in cylinders.

Solid Carbon Dioxide (Dry Ice): Carbon dioxide is a prime example. Under high pressure, CO₂ can be stored as a solid, known as dry ice. If the pressure is decreased to 1 atmosphere, solid CO₂ converts directly into gaseous CO₂ without becoming a liquid. This makes dry ice a useful cooling agent.

Thus, we can conclude that pressure and temperature determine the state of a substance.

Interconversion of the Three States of Matter:

• Solid to Liquid: Melting/Fusion
• Liquid to Gas: Vaporization/Boiling
• Gas to Liquid: Condensation
• Liquid to Solid: Solidification/Freezing
• Solid to Gas: Sublimation
• Gas to Solid: Deposition

Evaporation

The phenomenon of change of a liquid into vapours at any temperature below its boiling point is called evaporation. Unlike boiling, which is a bulk phenomenon (the entire liquid boils at once), evaporation is a surface phenomenon. Particles on the surface of the liquid with higher kinetic energy can break away from the forces of attraction and escape into the air as vapour.

Factors Affecting Evaporation

The rate of evaporation is influenced by several factors:

• Increase in Surface Area: Evaporation is a surface phenomenon. If the surface area is increased, the rate of evaporation increases. This is why we spread out wet clothes to dry them faster.
• Increase in Temperature: With an increase in temperature, more particles get enough kinetic energy to go into the vapour state. Therefore, evaporation is faster on a hot day than on a cold day.
• Decrease in Humidity: Humidity is the amount of water vapour present in the air. The air around us can hold only a definite amount of water vapour at a given temperature. If the amount of water in the air is already high (high humidity), the rate of evaporation decreases. Clothes dry slower on a humid day.
• Increase in Wind Speed: With an increase in wind speed, the particles of water vapour move away with the wind, decreasing the amount of water vapour in the immediate surroundings. This increases the rate of evaporation. We often see clothes drying faster on a windy day.

How Does Evaporation Cause Cooling?

The process of evaporation requires energy. The particles of the liquid absorb this energy from their surroundings to escape into the vapour phase. This absorption of energy from the surroundings makes the surroundings cool.

Real-life Examples of Cooling due to Evaporation:

• Acetone on Palm: When you put some acetone (nail polish remover) on your palm, the particles of acetone absorb energy from your palm and the surroundings and evaporate. This leaves your palm feeling cool.
• Sweating: During hot summer days, our bodies sweat. The sweat absorbs heat from our body surface and evaporates, which helps to keep our body cool. This is our body's natural cooling mechanism.
• Earthen Pot (Matka): People often store water in earthen pots during summer. An earthen pot has a large number of tiny pores on its surface. Water seeps out through these pores and evaporates from the surface of the pot. The energy required for this evaporation is taken from the water inside the pot, thus making the water cool.
• Wearing Cotton Clothes in Summer: We prefer cotton clothes in summer because cotton is a good absorber of water. It absorbs the sweat from our body and exposes it to the atmosphere for easy evaporation, which helps in cooling our body.

Important Questions and Answers

Question 1: Convert the following temperatures to the Celsius scale. (a) 300 K (b) 573 K

Answer: The relationship between the Kelvin (K) and Celsius (°C) scales is given by the formula: Temperature in Kelvin = Temperature in Celsius + 273.15. For simplicity in school curriculum, we often use 273.

So, Temperature in Celsius (°C) = Temperature in Kelvin (K) - 273.

(a) 300 K:

Temperature in °C = 300 - 273 = 27°C.

Therefore, 300 K is equal to 27°C.

(b) 573 K:

Temperature in °C = 573 - 273 = 300°C.

Therefore, 573 K is equal to 300°C.

Question 2: Why does a desert cooler cool better on a hot dry day?

Answer: A desert cooler works on the principle of cooling by evaporation. It has pads on its sides which are kept wet by a pump. A fan pulls hot, dry air from outside through these wet pads. As the air passes through the pads, the water on the pads evaporates.

The process of evaporation requires energy, which is taken from the hot air passing through. This absorption of heat causes the air to cool down significantly. The fan then blows this cool air into the room.

This process is most effective on a 'hot dry day' for two reasons:

• Hot Day (High Temperature): A higher temperature increases the rate of evaporation. More water evaporates, which means more heat is absorbed from the air, resulting in greater cooling.
• Dry Day (Low Humidity): A dry day means there is very little water vapour in the air (low humidity). This allows for a much higher rate of evaporation, as the air can easily accommodate more water vapour. On a humid day, the air is already saturated with moisture, so evaporation is slow and the cooling effect is greatly reduced.

Therefore, the combination of high temperature and low humidity on a hot dry day provides the ideal conditions for a desert cooler to work at maximum efficiency.

Question 3: Give two reasons to justify that— (a) water at room temperature is a liquid. (b) an iron almirah is a solid at room temperature.

Answer:

(a) Water at room temperature is a liquid because:

1. It has no fixed shape but has a fixed volume: Water takes the shape of the container it is poured into, a characteristic property of liquids. However, its volume remains constant. If you pour 1 litre of water from a bottle into a jug, it will still be 1 litre of water.
2. It can flow (fluidity): Water particles are not held in fixed positions and can slide past one another. This allows water to flow from a higher level to a lower level, which is a defining characteristic of liquids.

(b) An iron almirah is a solid at room temperature because:

1. It has a definite shape and a definite volume: An iron almirah maintains its rectangular shape and fixed volume regardless of where it is placed. It does not take the shape of the room. This rigidity is a hallmark of the solid state.
2. It is incompressible and rigid: The particles in the iron almirah are very tightly packed with strong intermolecular forces. This makes it rigid and almost impossible to compress under normal conditions.

Question 4: Why is ice at 273 K more effective in cooling than water at the same temperature?

Answer: Ice at 273 K (0°C) is more effective in cooling than water at 273 K (0°C) due to the concept of latent heat of fusion.

When you use ice for cooling, it absorbs heat from the surroundings in two stages:

1. First, the ice needs to melt into water. To do this, it must absorb a significant amount of energy from the surroundings to break the bonds of its solid structure. This absorbed energy is the 'latent heat of fusion' and it does not cause any rise in temperature. The temperature of the ice and the resulting water remains at 0°C until all the ice has melted. For every kilogram of ice that melts, it absorbs approximately 3.34 x 10^5 Joules of heat energy.
2. Once all the ice has melted into water at 0°C, this water will then start to absorb more heat, causing its temperature to rise above 0°C.

On the other hand, if you start with water at 0°C, it can only absorb heat to raise its temperature (stage 2). It does not have the initial melting phase (stage 1).

Therefore, ice at 273 K can \textract an additional amount of heat (the latent heat of fusion) from a substance compared to water at the same temperature. This makes ice a much more powerful cooling agent.

Chapter Summary

Here are the key takeaways from 'Matter in Our Surroundings':

• Matter is anything that has mass and occupies space.
• Matter is made up of \textremely small, discrete particles.
• The particles of matter have spaces between them, are in continuous random motion, and attract each other.
• The three main states of matter are solid, liquid, and gas, which are determined by the strength of interparticle forces and the kinetic energy of the particles.
• Solids have a definite shape and volume, are rigid, and are incompressible.
• Liquids have a definite volume but no definite shape, are fluid, and have low compressibility.
• Gases have no definite shape or volume, are highly fluid, and are highly compressible.
• The states of matter are interconvertible by changing temperature or pressure.
• Melting (Fusion): The process of a solid changing into a liquid. The temperature at which this occurs is the melting point.
• Boiling (Vaporization): The process of a liquid changing into a gas. The temperature at which this occurs is the boiling point.
• Latent Heat: The hidden heat energy absorbed or released during a change of state without a change in temperature (e.g., latent heat of fusion, latent heat of vaporization).
• Sublimation: The direct conversion of a solid to a gas, or vice versa (deposition), without passing through the liquid state.
• Evaporation: A surface phenomenon where a liquid turns into vapour at any temperature below its boiling point. It causes a cooling effect.
• The rate of evaporation depends on surface area, temperature, humidity, and wind speed.