Matter in our surroundings Class 9 Notes Science Chapter 1

When we observe our surroundings, we notice a vast array of objects, each differing in shape, size, and texture. Despite these differences, all these objects share a fundamental characteristic: they are made up of matter. But what exactly is matter?

Matter is defined as anything that occupies space and has mass. It is the substance that constitutes the entire universe, from the smallest grain of sand to the largest star in the sky. 

• Everything around us, whether it's the air we breathe, the water we drink, or the objects we use daily, is composed of matter.
• Universal Presence: Everything, from solids to gases, is made of matter.
• Historical and Modern Views: Early thinkers and modern scientists both sought to classify matter, though their methods differed.
• Focus of the Chapter: This chapter explores the physical properties of matter and its various states—solid, liquid, and gas.

By understanding these concepts, you'll gain insight into the material world that surrounds us.

Physical Nature of Matter

The physical nature of matter refers to its fundamental properties and behavior, as observed and studied through scientific inquiry. Let's see those properties.

1. Matter is made up of Particles

Matter is  made up of particles. These particles can be atoms, molecules, or ions, depending on the specific substance.

The particle nature of matter can be demonstrated by a simple activity.

• Experiment:
  (i) Take about 50 ml water in a 100 ml beaker.
  (ii) Mark the level of water.
  (iii) Add some salt to the beaker and stir with the help of a glass rod.
  (iv) Observe the change in water level.
  

• Observation:
  (i) It is observed that the crystals of salt disappear.
  (ii) The level of water remains unchanged.
• Explanation:
  A water molecule consists of hydrogen and oxygen atom, between hydrogen and oxygen, there are large empty spaces. These empty spaces are known as void.

Void: When we add salt to the water, it goes into that void. As a result, we do not see any change in volume.

• Conclusion: 
  This activity shows that matters are made of small particles. And there is space between these particles.

2. How Small are these Particles of Matter?

The size of particles of matter can vary widely depending on what type of particle you're considering and the scale at which you're measuring. Let's perform an experiment.

• Experiment:
  (i) Take a 250 ml beaker and add 100 ml water to it.
  (ii) Now add 2-3 crystals of potassium permanganate (KMnO₄) and stir with a glass rod in order to dissolve the crystals.
  (iii) Take 10 ml of this solution and add to 100 ml of water taken in another beaker.
  (iv) Take 10 ml of this diluted solution and put into 100 ml of water taken in still another beaker.
  (v) Repeat this process 10 times observe the colour of the solution in the last beaker.
  
  Diffusion of Potassium Permanganate in Water
• Observation:
  (i) When we add potassium permanganate in water, the colour of water changes to pink.
  (ii) Dilution decreases the colour intensity of the solution.
• Explanation:
  (i) A small amount of Potassium permanganate contain millions of its molecules. When we dissolve potassium permanganate in water, its molecule spread uniformly in the solution and give a pink appearance.
  (ii) Dilution lowers the amount of the particles in a subsequent solution. As a result, we see a lower colour intensity.
• Conclusion:  This activity proves that matter is made up of tiny particles. 

Characteristics of Particles of Matter

The characteristics of particles of matter encompass a range of properties that describe their behaviour, structure, and interactions. Let's see those characteristics.

1. Particles of Matter have Space between Them

• There are small voids between every particle in a matter. 
• This characteristic is the concept behind the solubility of a substance in other substances.

Activity

• Aim: To demonstrate the space between particles of matter.
• Experiment:
  (i) Take a glass of water.
  (ii) Put a teaspoon of salt/sugar and mix them properly. 
• Observation: The water is still clear. 
• Explanation: This is because the particles of salt/sugar get into the interparticle spaces between the water particles. 
• Conclusion:
  (i) This proves that there are voids between particles of a substance.
  (ii) If you add more salt/sugar, it will dissolve until all the space between water particles get filled.

2. Particles of Matter are Continuously Moving

• If an incense stick (Agarbatti) is lighted and placed in one corner of a room, its pleasant smell spreads in the whole room quickly.
  Agarbatti
• It demonstrates that the particles of matter possess motion. When we light an incense stick, it produces some gases (vapour) having a pleasant smell.
• The particles of these gases due to motion spread in the entire room. As a result, we can observe the smell of the lighted incense stick from a long distance.
• This shows that Matters consist of small particles which are moving continuously. This means that particles of matter possess kinetic energy.

Activity 

• Aim: To demonstrate that the Kinetic Energy of particles increases with an increase in temperature.
• Experiment:
  (i) Take two beakers. To one beaker add 100 ml of cold water and to the other beaker add 100 ml of hot water.
  (ii) Now add some crystals of potassium permanganate or copper sulphate to both the beakers.

Kinetic energy: Kinetic energy of an object is the measure of the work an object can do by virtue of its motion.  The kinetic energy is 1/2 mv². To accelerate an object, we have to apply force. To apply force, we need to do work. When work is done on an object, energy is transferred, and the object moves with a new constant speed. We call the energy that is transferred kinetic energy, and it depends on the mass and speed achieved. 

• Observation: 
  It is observed that crystals in hot water diffuse and dissolves faster than in a beaker containing cold water.
• Conclusion:
  (i) All substances have some kinetic energy in it. When we heat a substance, its kinetic energy increases.
  (ii) Heating water results in an increase in its kinetic energy; as a result, we see that crystals dissolve in much lesser time.
  (iii) From these activities, it is observed that when two different forms of matter are brought into contact, they intermix spontaneously.
  (iv) This intermixing is possible due to the motion of the particles of matter and also due to the spaces between them. 

3. Particles of Matter Attract each other

• There are some forces of attraction between the particles of matter which bind them together. The force of attraction between the particles of the same substance is known as cohesion.
• The force of attraction (or cohesion) is different in the particles of different kinds of matter. In general, the force of attraction is maximum in the particles of solid matter and minimum in the particles of gaseous matter.

Activity

• Aim: To demonstrate the attractive forces between particles of matter.
• Experiment:
  (i) Take a piece of iron wire, a piece of chalk and a rubber band.
  (ii) Try to break them by hammering, cutting or stretching.
• Observation:
  (i) Hammering a piece of the iron nail does not break the nail but flatten its surface.
  (ii) Hammering chalk breaks the chalk and gives us powdered chalk.
  (iii) We can stretch the rubber band to a large length without any break.
• Conclusion:
  (i) Since energy is required to break crystals of matter into particles.
  (ii) It indicates that particles in the matter are held by some attractive forces; the strength of these attractive forces varies from one matter to another.

States of Matter

The three states of matter are the distinct physical forms that matter can take: solid, liquid, and gas.

• Matter can exist in one of three main states: solid, liquid, or gas.
• Solid matter is composed of tightly packed particles. A solid will retain its shape; the particles are not free to move around.
• Liquid matter is made of more loosely packed particles. It will take the shape of its container. Particles can move about within a liquid, but they are packed densely enough that volume is maintained.
• Gaseous matter is composed of particles packed so loosely that it has neither a defined shape nor a defined volume. A gas can be compressed.

Three states of matter

Fig: System of Enthalpy

This diagram shows the nomenclature for the different phase transitions.

The Solid State

• A solid particles are packed closely together.
• The forces between the particles are strong enough that the particles cannot move freely; they can only vibrate. 
• As a result, a solid has a stable, definite shape and a definite volume. Solids can only change shape under force, as when broken or cut.

Fig: Structure of Solids

• In crystalline solids, particles are packed in a regularly ordered, repeating pattern. 
• There are many different crystal structures, and the same substance can have more than one structure.
• Example: Iron has a body-centered cubic structure at temperatures below 912°C and a face-centered cubic structure between 912 and 1394°C. Ice has fifteen known crystal structures, each of which exists at a different temperature and pressure.

The Liquid State

A liquid is a fluid that conforms to the shape of its container but that retains a nearly constant volume independent of pressure. 

• The volume is definite (does not change) if the temperature and pressure are constant. 
• When a solid is heated above its melting point, it becomes liquid because the pressure is higher than the triple point of the substance.
• Intermolecular (or interatomic or interionic) forces are still important, but the molecules have enough energy to move around, which makes the structure mobile. 
• This means that a liquid is not definite in shape but rather conforms to the shape of its container. 
• Its volume is usually greater than that of its corresponding solid (water is a well-known exception to this rule). 
• The highest temperature at which a particular liquid can exist is called its critical temperature.

Fig: Structure of Liquid

The Gaseous State

Gas molecules have either very weak bonds or no bonds at all, so they can move freely and quickly. 

• Because of this, not only will a gas conform to the shape of its container, it will also expand to completely fill the container. 
• Gas molecules have enough kinetic energy that the effect of intermolecular forces is small (or zero, for an ideal gas), and they are spaced very far apart from each other; the typical distance between neighboring molecules is much greater than the size of the molecules themselves. 
• A gas at a temperature below its critical temperature can also be called a vapor.
• A vapor can be liquefied through compression without cooling. It can also exist in equilibrium with a liquid (or solid), in which case the gas pressure equals the vapor pressure of the liquid (or solid).Fig: Structure of Gas

A supercritical fluid (SCF) is a gas whose temperature and pressure are greater than the critical temperature and critical pressure. In this state, the distinction between liquid and gas disappears. A supercritical fluid has the physical properties of a gas, but its high density lends it the properties of a solvent in some cases. This can be useful in several applications.
Example: Supercritical carbon dioxide is used to extract caffeine in the manufacturing of decaffeinated coffee.

Can Matter Change its State?

• A substance may exist in three states of matter i.e., solid, liquid or gas, depending upon the conditions of temperature and pressure. 
• By changing the conditions of temperature and pressure, all three states could be obtained (solid, liquid, gas). On heating, solid changes into a liquid which on further heating changes into a gas.
  Example: Water exists in all three states.
  Solid: Ice, Liquid: Water, Gas: Water Vapor.
• Ice is a solid state and may be melted to form water (liquid) which on further heating changes into steam (gas). These changes can also be reversed on cooling.

Note:

1. Changing a solid to a liquid is called melting.

2. Changing a liquid to solid is called solidification.

3. Changing a liquid to gas is called vaporization.

4. Changing a gas to liquid is called condensation.

5. Changing a solid to gas directly is called sublimation.

Temperature and pressure are the two factors which decide whether a given substance would be in a solid, liquid or gaseous state.

1. Effect of Change of Temperature

Activity

The effect of temperature on three states of matter could be seen by performing the following activity.

• Procedure
  (i) Take a piece of about 100 - 150 g of ice in a beaker.
  (ii) Hang a thermometer in it so that its bulb is in contact with ice.
  (iii) Start heating the beaker slowly on a low flame.
  (iv) Note down the temperature when ice starts changing of water & ice has been converted to water.
  (v) Record all observations for the conversion of solid ice into liquid water.
  (vi) Now, place a glass rod in the beaker and slowly heat the beaker with constant stirring with help of a glass rod.
  (vii) Note the temperature when water starts changing into water vapours.
  (viii) Record all observations for the conversion of water in the liquid state to vapour state.
• Observation: It is observed that as temperature increases, the ice starts changing into water. This change is called "Melting". The temperature remains same till all the ice changes into water. The thermometer shows 0°C until all the ice has melted. On further heating, the temperature starts rising. At 373 K (or 100°C), water starts boiling. As the water continues to boil, the temperature remains almost constant.
• Conclusion of the above activity 
  This experiment demonstrates that we can change the physical state of matter by heating (Solid → Liquid → Gas).

Solid to Liquid Change (Melting) 

• Ice is solid. In solids, the particles are tightly packed together. When we heat a solid, its particles become more energetic and kinetic energy of the particles increases. Due to the increase in kinetic energy, the particles start vibrating more strongly with greater speed. 
• The energy supplied by heat overcomes the intermolecular forces of attraction between the particles. As a result, the particles leave their mean position and break away from each other. After this, solid melts and a liquid is formed.
• The temperature at which solid melts to become a liquid at atmospheric pressure is called its Melting point(or fusion). 
• The process of melting is also called "Fusion". The melting point of ice is 0°C. It may also be written as 273.16 K or 273 K.
• The temperature is represented in Celsius scale as °C. Nowadays, it is expressed on Kelvin scale (K).
• Conversion of temperature on Celsius scale to the Kelvin scale
  Example:
  0°C = 0 + 273 = 273 K
  100°C = 100 + 273 = 373 K
• Conversion of temperature on Kelvin scale to Celsius scale
  Example:
  -373 K = 373 - 273 = 100°C
  273 K = (273 - 273) = 0°C

Liquid to Gas Change (Boiling or Vapourization)

• When we continue heating the water, at 100°C water molecules again break the force of attraction and become vapour. Here, 100°C acts as an equilibrium point, and any energy beyond this is utilized to break the force of attraction. So the temperature also remains the same, i.e. 100°C. We call this latent heat as latent heat of evaporation.
• The temperature at which a liquid changes into a gas or vapor at the atmospheric pressure is called its boiling point.
  Example: For water, the boiling point is 100°C or 373 K. The particles in steam i.e., water vapour at 373 K have more energy than water at the same temperature. Reason: This is because the particle in steam has absorbed extra energy in the form of latent heat of vaporization.
• The boiling point of a liquid also indicates the strength of the intermolecular force of attraction between particles. Volatile liquids such as alcohol, petrol, and acetone have very weak intermolecular forces. Therefore, they boil at a low temperature. On the other hand, water has stronger inter-molecular forces of attraction and therefore, it boils at a higher temperature. When steam is cooled, it condenses to water & when water is cooled, it changes to ice.
  Example: If the boiling point of water is 100°C (373 K), then the liquefaction point of steam is 100°C (373 K).

2. Effect of Change of Pressure

• Gases are compressible because of applying pressure, the space between the gaseous particles decreases. Therefore gases can be compressed easily.
• By applying pressure and reducing temperature, the gases can be converted into liquids i.e. gases will be liquefied.
• This process of conversion of a gas into a liquid by increasing pressure or decreasing temperature is called liquefaction.
• Solid carbon dioxide is also called dry ice. Solid CO2 gets converted into gaseous state directly on decreasing pressure to 1 atmosphere without coming into the liquid state.

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

Latent Heat (Hidden Heat)

• It is observed that the temperature of the system does not change after melting point is achieved until all the ice melts, though we continue to heat the beaker. 
• This is happening because the heat supplied is used up in changing the state by breaking the intermolecular forces of attraction which holds them in the solid state. As a result, there is no change in temperature until all the ice melts. This energy required to change a solid into liquid is called "latent heat". The word "latent" means "hidden" because this energy is hidden into the contents of the beaker.
• Latent heat is of two types:
  (i) Latent heat of fusion.
  (ii) Latent heat of vaporization.

Latent Heat of Fusion 

• Latent heat of fusion is defined as the amount of heat energy required to change 1 kg of a solid into a liquid at atmospheric pressure without any change in temperature at its melting point. 
• The latent heat of fusion of ice is 3.34 × 105 J/Kg.
• The numerical value of melting point and freezing point is the same.
  Example: If the melting point of ice is 0°C (273 K), then the freezing point of water is 0°C (273 K).

Latent Heat of Vaporization 

• The latent heat of vaporization of a liquid is the quantity of heat in joules required to convert 1 kilogram of the liquid (at its boiling point) to vapour or gas, without any change in temperature. 
• The latent heat of vaporization of water is 22.5 × 10⁵ joules per kilogram (or 22.5 × 10⁵ J/kg).

Condensation

Formation of Water droplets on the surface of a cold bottle due to the condensation of water vapour present in the air.

• The process of changing a gas (or vapour) to a liquid by cooling is called condensation where gas is cooled enough.
• So, when steam (or water vapour) changes into water on cooling, it is called condensation of steam (or condensation of water vapour).
• It is the reverse of vaporization. (Boiling)

Freezing

• The process of changing a liquid (solidification) into a solid by cooling is called freezing. 
• When a liquid is cooled, its particles lose energy due to which they move slowly.
• If the liquid is cooled enough (upto freezing point) its each particle stops moving and vibrates about a fixed position. At this stage, the liquid freezes and becomes a solid.
• Freezing is the reverse of melting.

Sublimation

Sublimation

• The process due to which a solid directly changes into gaseous state on heating, without changing first into liquid state and the gaseous state, directly changes into solid state on cooling, is known as "Sublimation".
  Example: Ammonium chloride, Camphor, Iodine, Naphthalene, Solid carbon dioxide or (dry Ice).
• Sublime: A gaseous form, directly formed from a solid on heating, is known as sublime.
• Sublimate: A solid state of matter formed directly from its gaseous state on cooling is called sublimate.
• To understand sublimation process we can do an activity:
  > Take some camphor or ammonium chloride.
  > Powder it and put in a china dish.
  > Place an inverted funnel over the china dish.
  > Put a cotton plug on the stem of the funnel, as shown in the Figure below.
  > Heat the china dish slowly.
  
  Sublimation
• Observation: Heating camphor or ammonium chloride results in its vapour formation. These vapour cool at the upper part of the beaker and deposit there.
• Explanation: Solid state is directly converted into the gaseous state. This experiment shows the sublimation process. Cotton plug on the beaker prevents the escape of the formed vapour. In the upper part of the beaker, the temperature is lesser. As a result, the vapours cool down and deposit as solids on the neck of the beaker. 
• This process of conversion of gas to solid is inverse of sublimation and we call this process Deposition.

Evaporation

• The process of a liquid changing into vapour (or gas) even below its boiling point is called evaporation.
• Evaporation of a liquid can take place even at room temperature, though it is faster at higher temperatures. It is the surface phenomenon because it occurs at the surface of a liquid only. Whatever be the temperature at which evaporation takes place, the latent heat of vaporization must be supplied whenever a liquid changes into vapour (or gas).

Explanation about Evaporation 

• Some particles in a liquid always have more kinetic energy than the others. So, even when a liquid is well below its boiling point, some of its particles have enough energy to break the forces of attraction between the particles and escape from the surface of the liquid in the form of vapour (or gas). 
• Thus, the fast-moving particles (or molecules) of a liquid are constantly escaping from the liquid to form vapour (or gas).
• Examples:
  (i) Water in ponds changes from liquid to vapour without reaching the boiling point.
  (ii) Water when left uncovered slowly changes into vapours.
  (iii) When we put wet clothes for drying, the water from the clothes goes to the atmosphere.

Differences between Evaporation and Boiling

Factors Affecting Evaporation

There are five factors which affect the rate of evaporation:

• Nature of liquid: Different liquids have different rates of evaporation. A liquid having weaker interparticle attractive forces evaporates at a faster rate because less energy is required to overcome the attractive forces.
  Example: Acetone evaporates faster than water.
• The surface area of the liquid: The evaporation depends upon the surface area. If the surface area is increased, the rate of evaporation increases because the high energy particles from the liquid can go into the gas phase only through the surface.
  Example:
  (i) The rate of evaporation increases when we put kerosene or petrol in an open china dish than in a test tube.
  (ii) Clothes dry faster when they are well spread because the surface area for evaporation increases.
• Temperature: Rate of evaporation increases with increase in temperature. This is because with the increase in temperature number of particles get enough kinetic energy to go into the vapour state (or gaseous state).
  Example: Clothes dry faster in summers than in winters.
• Humidity in the air: The air around us contains water vapour or moisture. The amount of water present in the air is referred to as humidity. The air cannot hold more than a definite amount of water vapour at a given temperature. If the humidity is more, the rate of vaporization decreases. The rate of evaporation is more if the air is dry.
  Example: Clothes do not dry easily during the rainy season because the rate of evaporation is less due to high moisture content (humidity) in the air.
• Wind speed: The rate of evaporation also increases with increase in speed of the wind. This is because with increase in speed of wind, the particles of water vapour move away with wind resulting in a decrease in the amount of vapour in the atmosphere.
  Example:
  (i) Clothes dry faster on a windy day.
  (ii) In a desert cooler an exhaust fan sucks the moist air from the cooler chamber which results in a greater rate of evaporation of water and hence greater cooling.

How does Evaporation cause Cooling?

• During evaporation, cooling is always caused. This is because evaporation is a phenomenon in which only the high energy particles leave the liquid surface. As a result, the particles having low energy are left behind. Therefore, the average molecular energy of the remaining particles left in the liquid state is lowered. As a result, there is a decrease in temperature on the part of the liquid that is left. Thus evaporation causes cooling.
• Examples:
  (i) When we pour some acetone on our palm, we feel cold. This is because the particles gain energy from our palm or surroundings and leave the palm feeling cool.
  (ii) We sprinkle water on the root of the open ground after a sunny hot day. This cools the roof or open ground. This is because the large latent heat of vaporization of water helps to cool the hot surface.

Some other examples of Evaporation

• We should wear cotton clothes in hot summer days to stay cool and comfortable: This can be explained as follows. We get a lot of sweat on our bodies on hot summer days. Cotton is a good absorber of water, so it absorbs the sweat from our body and exposes it to the air for evaporation. The evaporation of this sweat cools our body. Synthetic clothes (made of polyester etc.) do not absorb a lot of sweat, so they fail to keep our body cool in summer.
• We see water droplets on the outer surface of a glass containing ice-cold water: Take some ice-cold water in a glass. Soon we will see water droplets on the outer surface of the glass.
• The water vapour present in the air, on coming in contact with the cold glass of water loses energy and gets converted to the liquid state, which we see as water droplets.
• Water keeps cool in the earthen pot (matki) during summer: When the water oozes out of the pores of an earthen pot, during hot summer, it evaporates rapidly. As the cooling is caused by evaporation, therefore, the temperature of water within the pot falls and hence it becomes cool.

Earthen Pot

• Rapid cooling of hot tea: If tea is too hot to sip, we pour it into the saucer. In doing so, we increase the surface area and the rate of evaporation. This, in turn, causes cooling and the tea attains the desired temperature for sipping.
• A wet handkerchief is placed on the forehead of a person suffering from a high fever: The logic behind placing wet cloth is that as the water from the wet cloth evaporates, it takes heat from the skull and the brain within it. This, in turn, lowers the temperature of the brain and protects it from any damage due to high temperature.
• We often sprinkle water on the road in summer: The water evaporates rapidly from the hot surface of the road, thereby taking heat away from it. Thus, the road becomes cool.