Introduction

Atoms and molecules are the fundamental building blocks of matter. A clear understanding of atoms and molecules explains why substances combine in particular ways, why physical and chemical properties differ from one substance to another, and how new substances are formed in chemical reactions.

• Maharishi Kanad and Pakudha Katyayama in ancient India proposed that matter can be divided into smaller indivisible particles called Parmanu.
• Democritus and Leucippus in ancient Greece proposed a similar idea: matter is made of indivisible particles called atoms.
• These early ideas were philosophical and lacked experimental proof until modern chemistry developed in the 18th century.
• In the late 18th century, Antoine L. Lavoisier established quantitative methods in chemistry and laid foundations for modern chemical science by formulating laws about chemical combinations.
• Lavoisier and Joseph L. Proust performed careful experiments that led to two important laws: the Law of Conservation of Mass and the Law of Constant Proportions (also called the Law of Definite Proportions).
• These laws guided later work and helped John Dalton formulate his atomic theory that explained why these laws hold true for chemical reactions.

Try yourself:

Who proposed the idea that matter can be divided into smaller particles called Parmanu?

• A.Maharishi Kanad
• B.Democritus
• C.Antoine L. Lavoisier
• D.Joseph L. Proust

Laws of Chemical Combination

Two fundamental laws describe how substances combine in chemical reactions: the Law of Conservation of Mass and the Law of Constant Proportions. These laws provide the experimental basis for the atomic view of matter.

1. Law of Conservation of Mass

The Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.

Example of Law of conservation of Mass

Careful experiments – for example, mixing chemical solutions in closed containers and measuring mass before and after reaction – show that the measured total mass remains unchanged, supporting this law.

Try yourself:

According to the Law of Conservation of Mass, what happens to the mass during a chemical reaction?

• A.The mass increases.
• B.The mass decreases.
• C.The mass remains constant.
• D.The mass is converted into energy.

2. Law of Constant Proportion

The Law of Constant Proportion (Law of Definite Proportions) states that a chemical compound always contains the same elements in the same fixed proportion by mass, irrespective of its source.

Example of Law of Constant Proportion

For example, pure water always contains hydrogen and oxygen in the mass ratio 1 : 8. This proportion is the same whether the water comes from a river, a well, or rain.

John Dalton’s Atomic Theory

To explain these laws, John Dalton proposed an atomic theory which gave an experimental and conceptual basis for atoms and compounds.

John Dalton

Postulates of Dalton’s Atomic Theory

1. All matter is made of extremely small particles called atoms.
2. Atoms are indivisible by chemical means and remain unchanged in chemical reactions.
3. Atoms of the same element are identical in mass and properties.
4. Atoms of different elements have different masses and properties.
5. Atoms combine in simple whole-number ratios to form compounds.
6. The relative number and types of atoms in a given compound are constant (fixed composition).

Background on John Dalton

• John Dalton was born in 1766. His atomic hypothesis explained the Law of Conservation of Mass and the Law of Definite Proportions quantitatively and conceptually.

What is an Atom?

An atom is the smallest particle of an element that retains the chemical properties of that element and cannot be broken down by chemical means.

• Atoms are extremely small; a very large number of atoms are required to form visible matter.
• A layer only a few million atoms thick may be comparable in thickness to a sheet of paper.

Atomic Radius

The atomic radius is a measure of the size of an atom, typically expressed in nanometres (nm), where 1 nm = 10^-9 m.

Try yourself:

Which statement best describes the Law of Constant Proportion?

• A.It states that in a chemical substance, elements are always present in definite proportions by volume.
• B.It states that in a chemical substance, elements are always present in definite proportions by mass.
• C.It states that in a chemical substance, elements are always present in indefinite proportions by mass.
• D.It states that in a chemical substance, elements can be present in any proportions by mass.

Modern Symbols of Elements

Element symbols evolved from early pictorial symbols to the simple one- or two-letter symbols used today. The International Union of Pure and Applied Chemistry (IUPAC) standardises these symbols.

• Historical background: John Dalton first used symbols to represent atoms. Later, Berzelius proposed using one or two letters derived from the element name to represent elements.
• Origin of element names: Some names come from places (e.g., copper from Cyprus) or from colours and other properties.
• Modern symbols: Most element symbols are derived from their English names; the first letter is capitalised and a second letter, if used, is lowercase.
• First letter + another letter: Examples include Chlorine: Cl, Zinc: Zn.
• Names from other languages: Some symbols come from Latin, Greek or other languages: Iron: Fe (ferrum), Sodium: Na (natrium), Potassium: K (kalium).

Symbols for Some Elements

Try yourself:

Which scientist pioneered the use of symbols for elements?

• A.Berzelius
• B.Dalton
• C.IUPAC
• D.None of the above

Atomic Mass

Atomic mass of an atom is the mass of that atom expressed relative to a standard. The standard used internationally is defined as 1/12 of the mass of a carbon-12 atom. The unit for atomic mass is the unified atomic mass unit (symbol u, also called amu).

• Atomic mass is the combined mass of protons, neutrons and electrons in an atom; in practice the electron mass is very small relative to protons and neutrons and often neglected in simple calculations.
• Atomic masses reported on the periodic table are average values that reflect the natural isotopic composition of the element.

Atomic mass of some elements

How Do Atoms Exist?

• Many atoms do not exist freely under normal conditions; they combine to form molecules or form ions that aggregate into ionic structures.
• Visible matter is composed of huge numbers of molecules or ionic units assembled together.

What is a Molecule?

A molecule is the smallest particle of an element or compound that can exist independently and retain the chemical properties of that substance.

Molecules of Elements

1. Monoatomic molecules: Some elements exist as single atoms (monoatomic) in their natural gaseous state, e.g., Helium (He), Argon (Ar).
2. Diatomic molecules: Several non-metals exist as molecules of two atoms, e.g., Hydrogen (H₂), Oxygen (O₂), Nitrogen (N₂), Chlorine (Cl₂).
3. Polyatomic molecules: Some elements form molecules with more than two atoms, e.g., Phosphorus (P₄), Sulphur (S₈).

Try yourself:What is the atomic mass of an atom?

• A. The total mass of the neutrons and protons in an atom.
• B.The mass of a carbon-12 atom in its ground state.
• C.The average mass of a group of atoms.
• D.The mass of an atomic particle.

Atomicity

Atomicity is the number of atoms present in one molecule of an element.

Atomicity of some non-metals

Molecules of Compounds

When atoms of different elements combine, they form molecules of compounds. These molecules have fixed compositions and properties different from their constituent elements.

What is an Ion?

• An ion is an atom or a group of atoms that carries an electric charge due to loss or gain of electrons.
• A positively charged ion is called a cation; a negatively charged ion is called an anion.
• Compounds formed from metals and non-metals often contain ions; such compounds are called ionic compounds.
• Polyatomic ions are groups of atoms bonded together that carry a net charge, for example, NO₃^–, SO₄^2-, OH^–.

Writing Chemical Formulae

• The chemical formula of a compound shows which elements are present and the number of atoms of each element in the smallest unit of that compound.
• To write formulae, you must know element symbols and the valencies (combining capacities) or ionic charges of the atoms/ions involved.
• Valency indicates how many electrons an atom gains, loses or shares when it forms a compound.
• Think of valency as the number of bonds an atom typically forms: it is the atom’s “combining power”.

Rules for Formula Writing

• The total positive charge and total negative charge in a neutral compound must balance.
• When writing formulae for compounds of a metal and a non-metal, write the metal first and the non-metal second (e.g., CaO, NaCl).
• Use the simplest whole-number ratio of atoms or ions that balances charges.
• When polyatomic ions are present in more than one number, enclose the polyatomic ion in brackets and write the number outside the brackets, e.g., Mg(OH)₂.
• Practice with examples to become familiar with common valencies and formulas.

Try yourself:What is the atomicity of a molecule?

• A.The number of atoms in a molecule
• B.The number of ions in a molecule
• C.The number of elements in a molecule
• D.The number of protons in a molecule

Formulae of Simple Compounds

Binary compounds (formed by two elements) can be written by criss-crossing valencies or balancing charges of ions.

Example:

• Carbon tetrachloride, CCl₄: Carbon (valency 4) combines with chlorine (valency 1) to give the formula CCl₄.

• Magnesium chloride, MgCl₂: Magnesium (valency 2) combines with chlorine (valency 1) to give MgCl₂.

Molecular Mass

The molecular mass (relative molecular mass) of a molecule is the sum of the atomic masses of all atoms present in the molecule. It is expressed in atomic mass units (u).

Example 1:

(a) Calculate the relative molecular mass of water (H₂O).

(b) Calculate the molecular mass of HNO₃.

Solution:

(a)

Atomic mass of hydrogen = 1 u.

Atomic mass of oxygen = 16 u.

The molecular mass of H₂O = 2 × (atomic mass of H) + 1 × (atomic mass of O).

The molecular mass of H₂O = 2 × 1 + 16 = 18 u.

(b)

Atomic mass of hydrogen = 1 u.

Atomic mass of nitrogen = 14 u.

Atomic mass of oxygen = 16 u.

The molecular mass of HNO₃ = 1 × (atomic mass of H) + 1 × (atomic mass of N) + 3 × (atomic mass of O).

The molecular mass of HNO₃ = 1 + 14 + 3 × 16 = 63 u.

Try yourself:What is the chemical formula for magnesium chloride?

• A.MgC₂
• B.ZnCl₂
• C.MgCl₂
• D.Mg

Formula Unit Mass

Formula unit mass is the sum of the atomic masses of the atoms present in the formula unit of an ionic compound. It is calculated the same way as molecular mass but applied to ionic formula units.

Example 2: Calculate the formula unit mass of CaCl₂.

Solution:

The atomic mass of Ca = 40 u.

The atomic mass of Cl = 35.5 u.

The formula unit mass of CaCl₂ = atomic mass of Ca + 2 × atomic mass of Cl.

The formula unit mass of CaCl₂ = 40 + 2 × 35.5 = 40 + 71 = 111 u.

Introduction

Have you ever wondered if the water you drink, the air you breathe, or even the food you eat is completely pure? The chapter delves into the fascinating world of matter, where we explore how everything around us, from a simple sugar cube to the air we inhale, is made up of pure substances or mixtures. You’ll discover how mixtures can be separated into their components, and how pure substances are the building blocks of everything we see and use! 

What Is a Mixture?

Mixtures are constituted by more than one kind of pure form of matter, known as a substance.  For example, sea water, minerals, soil etc., are all mixtures.

•  When we say that something is pure, it means that all the constituent particles of that substance are the same in their chemical nature. A pure substance always consists of a single type of particle.
• When we look around, we can see that most of the matter around us exist as mixtures of two or more pure components.

Examples of Mixtures

• Dissolved sodium chloride can be separated from water by the physical process of evaporation but sodium chloride cannot be separated into sodium and chlorine by physical means.

Types of mixtures 

Depending upon the nature of components that form a mixture, we can have different types of mixtures. 

• Homogeneous mixtures: Mixtures which have a uniform composition throughout, are called homogeneous mixtures. For example, salt in water and sugar in water. 
• Heterogeneous mixtures: Mixtures which contain physically distinct parts and have non-uniform composition are called heterogeneous mixtures. For example, mixture of sodium chloride and iron filings, salt and sulphur.

Activity: Perform an activity to differentiate between solution, suspension and colloidal solution. 

Procedure:

• Distribute the following samples to four groups A, B, C and D of a class.
• A few crystals of copper sulphate to group A.
• One spatula is full of copper sulphate to group B.
• Chalk powder to group C.
• A few drops of milk or ink to group D.
• Ask each group to add the sample to water and stir using a glass rod.
• Direct a beam of light from a torch through the beakers.
• Leave the mixture undisturbed for a few minutes.
• Filter the mixtures

Solution, Suspension and Colloidal Solution

Observations: 

• We observe that groups A and B get a clear solution of copper sulphate although with different colour density.
• Group C get a suspension of chalk, which on filtration gives a residue of chalk on the filter paper and clear filtrate containing water.
• Group D get a colloidal solution of milk. The solution in this case is not transparent. But no suspension is obtained here and on filtration, no residue is obtained on the filter paper.

Conclusion:

• Group A and B created homogeneous mixtures with uniform composition, while Groups C and D made heterogeneous mixtures with distinct parts.
• This activity illustrates the difference between homogeneous and heterogeneous mixtures based on composition and appearance.

Try yourself:

What is a homogeneous mixture?

• A.A mixture that contains physically distinct parts and has a non-uniform composition.
• B.A mixture that has a uniform composition throughout.
• C.A mixture that can be separated into its individual components by physical means.
• D.A mixture that consists of a single type of particle.

What is a Solution?

A solution is a homogeneous mixture of two or more substances. Lemonade and soda water are example of solutions.

 A solution is not necessarily a liquid containing a solid, liquid or gas dissolved in it. Solid solution (alloys) and gaseous solution are also possible.

Alloys

Alloys are homogeneous mixtures of metals and cannot be separated into their components by physical methods. For example, brass is a mixture of approximately 30% zinc and 70% copper.

Solvent and solute

Solvent and solute are the components of the solution. 

• Solvent: The component that dissolves the other component in it is the solvent. 
• Solute: The component that is dissolved in the solvent is called solute.

For example, Tincture of iodine is a solution of iodine in alcohol. Aerated drinks like soda water are solutions of carbon dioxide as solute and water as solvent. Air is a mixture of a gas in a gas. The two major components of air are nitrogen (78%) and oxygen (21%).

Note: Generally solute is present in smaller quantity and solvent is present in greater quantity. For example, we have a solution of sugar in water in which case sugar is solute and water is the solvent.

Properties of a solution

• A solution is a homogeneous mixture.
• Particles of a solution are smaller than 1 nm (10^-9metre) in diameter. Therefore, they cannot be seen by naked eye.
• Because of small size, they do not scatter light.
• Solute particles cannot be separated from the mixture by filtration.

Concentration of the solution

Concentration of a solution is the amount of solute present in a given amount (mass or volume) of solution.

• A concentrated solution contains a large concentration of the solute in the solvent while a dilute solution contains a small concentration of the solute in the solvent.
• Mass by mass percentage of a solution
• Mass by volume percentage of a solution

Saturated solution 

At any particular temperature, a solution that has dissolved as much solute as it is capable of dissolving, is called saturated solution. 

No more solute can be dissolved in the saturated solution at a given temperature.

Solubility: The amount of solute present in a saturated solution at a given temperature is called its solubility

Unsaturated solution

If the amount of solute contained in a solution is less than saturation level, it is called unsaturated solution. 

Different substances in a given solvent have different solubilities at the same temperature.

What is a Suspension?

A suspension is a heterogeneous mixture in which the solute particles do not dissolve but remain suspended throughout the bulk of medium. 

• Particles of a suspension are visible to the naked eye. 
• For example, chalk powder in water.

Properties of a suspension

• Suspension is a heterogeneous mixture.
• Particles of suspension can be seen with a naked eye.
• Particles of a suspension scatter light passing through it and make its path visible.
• Solute particles in a suspension settle down after some time when kept undisturbed.
• Components of a suspension can be separated by the process of filtration.

What is a Colloidal solution?

A colloidal solution is a heterogeneous mixture in which the solute particles do not settle down but remain suspended. 

• Here the particle size of the solute is between 1 nm to 100 nm. 
• Colloidal particles cannot be seen with a naked eye but they scatter light thus making the path of light visible. 
• For example, milk and starch solution.

 Solution of Copper Sulphate does not show Tyndall Effect, Mixture of Water and Milk shows Tyndall Effect

Properties of colloidal solutions

•  A colloidal solution is a heterogeneous mixture.
• The particles of a colloid cannot be seen with a naked eye.
• Colloidal particles scatter light.
• Colloidal particles do not settle down when left undisturbed.
• Colloidal particles cannot be separated from the mixture by the process of filtration.

Dispersed phase and dispersion medium 

• These are the components of a colloidal solution. 
• The solute-like component in a colloidal solution are dispersed phase and the solvent like component in a colloidal solution is dispersed medium.

Some common examples of colloids

Physical and Chemical changes

A change which occurs without a change in composition and chemical nature of the substance is called physical change.

• Here a change only in physical properties of the substance takes place. 
• Properties like colour, hardness, rigidity, fluidity, density, melting point and boiling point are known as physical properties.
•  Melting of ice or boiling of water is a physical change because ice, water and water vapours are chemically the same substance i.e., H₂0.

A change of materials into another, new materials with different properties and one or more than one new substances are formed is called chemical change. 

• Burning is a chemical change. 
• During this process, one substance reacts with another substance to undergo a change in chemical composition. 
• During burning of candle, actually both physical and chemical changes take place. 
• The physical change involves the melting of wax and the chemical change involves the burning of wax into carbon dioxide and water.

Try yourself:

What is a solution?

• A.A heterogeneous mixture of two or more substances.
• B.A homogeneous mixture of two or more substances.
• C.A mixture of metals that cannot be separated by physical methods.
• D.A mixture in which solute particles do not dissolve but remain suspended.

What are the Types of Pure Substances?

On the basis of their chemical composition, substances can be classified either as elements or compounds.

Elements

• Lavoisier, a French chemist defined an element as the basic form of matter that cannot be broken down into simpler substances by chemical reactions. 
• Elements can be divided into the following main threetypes of substances:
  1. Metals.
  2. Non-metals.
  3. Metalloids.

Metals show the following properties

• They have a Lustre.
• They have silvery-grey or golden-yellow colour.
• They conduct heat and electricity.
• They are ductile that means they can be drawn into thin wires.
• They are malleable. That means they can be beaten into thin sheets.
• They are sonorous i.e., they make a ringing sound when hit.
• Examples of metals are gold, silver, copper, iron, sodium, etc.

Non-metals show the following properties

• They display a variety of colours.
• They are poor conductors of heat and electricity.
• They are not lustrous, sonorous or malleable.
• Examples of non-metals are oxygen, iodine, carbon, etc.
• Some elements have intermediate properties between those of metals and non-metals. 
• They are called metalloids. 
• Examples of metalloids are boron, silicon and germanium.

Some facts about elements

• The number of elements known at present is more than 100. Ninety two elements are naturally occurring and the rest are man-made.
• Majority of the elements are solids.
• Eleven elements are in gaseous state at room temperature.
• Two elements are liquid at room temperature – mercury and bromine.
• Elements gallium and cesium become liquid at a temperature slightly above room temperature (303 K).

Compounds

A compound is a substance composed of two or more elements chemically combined with one another in a fixed proportion.

Activity – Exploring the Properties of Iron and Sulphur Mixture.

Materials required: Crushed iron filings, sulphur, china dish, burner.

Procedure: 

• Divide the class into two groups.
• Provide each group with 5 g of iron filings and 3 g of sulphur powder in a china dish.

Group I:

• Mix and crush iron filings and sulphur powder together.

Group II:

• Mix and crush iron filings and sulphur powder together.
• Heat the mixture strongly until it becomes red hot.
• Remove from flame and let the mixture cool down.

Both Groups:

• Check for magnetism in the material obtained by bringing a magnet near it.
• Compare the texture and color of the material obtained by both groups.
• Add carbon disulphide to one part of the material obtained, stir well, and filter.
• Add dilute sulphuric acid or dilute hydrochloric acid to another part of the material obtained. (Note: Teacher supervision is necessary for this activity).
• Perform all the above steps with iron and sulphur separately.

Observation: Upon heating, iron and sulfur react chemically to form a compound. This compound has different properties from the original elements, indicating a chemical change. The mixture of iron and sulfur before heating shows the individual properties of both substances, but once heated, a new substance with distinct properties is created.

Conclusion

• When iron and sulfur are mixed and heated:
• Group I demonstrates a physical change, resulting in a mixture with similar properties to the individual substances (iron and sulfur).
• Group II exhibits a chemical change, where iron and sulfur react to form a compound with different properties.

This experiment highlights the differences between physical and chemical changes, as well as the concepts of mixtures and compounds in chemistry.

Mixture

If we simply mix iron filings with powdered sulphur and grind them together (no heating), we. obtain a mixture.

Comparison between mixtures and compounds

Table: Mixtures and Compounds

• We can summarise the physical and chemical nature of matter as under:

Try yourself:

What are the three main types of elements?

• A.Metals, non-metals, metalloids.
• B.Solids, liquids, gases.
• C.Compounds, mixtures, solutions.
• D.Protons, neutrons, electrons.

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 the 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, including the air we breathe, the food we eat, stones, clouds, stars, plants, animals, and even a tiny drop of water or a grain of sand—everything is matter.
• Universal Presence: Everything, from solids to gases, is made of matter.
• Historical Context: Ancient Indian philosophers identified matter as five basic elements: air, earth, fire, sky, and water, known as the ‘Panch Tatva’.
• 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 behaviour, as observed and studied through scientific inquiry. The main properties of matter are: 

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 of 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 atoms; between hydrogen and oxygen, there are large empty spaces. These empty spaces are known as voids. (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 matter is 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.

Procedure

(i) Take a 250 ml beaker and add 100 ml of water to it.

(ii) Now add 2-3 crystals of potassium permanganate (KMnO₄) and stir with a glass rod to dissolve the crystals.

(iii) Take 10 ml of this solution and add it to 100 ml of water taken in another beaker.

(iv) Take 10 ml of this diluted solution and put it into 100 ml of water taken in a still another beaker.

(v) Repeat this process 10 times observe the colour of the solution in the last beaker.

Observation:
(i) When we add potassium permanganate to water, the colour of the water changes to pink.
(ii) Dilution decreases the colour intensity of the solution.

Explanation:
(i) A small amount of Potassium permanganate contains millions of its molecules. When we dissolve potassium permanganate in water, its molecules 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.

Try yourself:

What happens to the color of water when potassium permanganate is added?

• A.It turns green
• B.It turns pink
• C.It turns blue
• D.It stays clear

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 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 is filled.

2. Particles of Matter are Continuously Moving

• If an incense stick (Agarbatti) is lit and placed in one corner of a room, its pleasant smell spreads throughout 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 throughout the entire room. As a result, we can observe the smell of the lit 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 dissolve faster than in a beaker containing cold water.

Conclusion:
(i) All substances have some kinetic energy. 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 a much shorter 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 flattens 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.

Try yourself:

What are particles of matter known to have?

• A.Color
• B.Light
• C.Mass
• D.Sound

States of Matter

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

Three States of Matter

• 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.

The Solid State

• 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-centred cubic structure at temperatures below 912°C and a face-centred 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 neighbouring 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 vapour 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 vapour 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.

Try yourself:

What shape does a solid retain?

• A.Liquid shape
• B.Variable shape
• C.Defined shape
• D.No shape

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 Vapour.
• 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:

– Changing a solid to a liquid is called melting.

– Changing a liquid to solid is called solidification.

– Changing a liquid to gas is called vaporization.

– Changing a gas to liquid is called condensation.

– Changing a solid to gas directly is called sublimation.
– Changing a gas to solid directly is called deposition.

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

Let’s Start with an 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 to 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 vapour.
(viii) Record all observations for the conversion of water in the liquid state to the vapour state.

Observation: It is observed that as the temperature increases, the ice starts changing into water. This change is called “Melting“. The temperature remains the 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).

​

(a) Melting of Ice

• When we increase the temperature of a solid, the kinetic energy of its particles also increases. This is because the particles start to vibrate more quickly. As the kinetic energy increases, it becomes strong enough to overcome the forces of attraction holding the particles together in fixed positions.
• Eventually, the particles are able to break free from their fixed positions and start moving more freely, leading to the melting of the solid into a liquid. The melting point is the minimum temperature at which this happens, and it varies depending on the strength of the forces of attraction between the particles.
• For example, the melting point of ice is 273.15 K (or 0°C ). The process of melting, also known as fusion, occurs when a solid is heated to its melting point.
• During the melting process, the temperature of the substance remains constant, even though heat is still being supplied. This is because the heat energy is being used to overcome the forces of attraction between the particles, rather than increasing the temperature.
• The heat energy required to change 1 kg of a solid into a liquid at its melting point is called the latent heat of fusion. This energy is absorbed by the particles without causing a rise in temperature, which is why it is called latent, meaning hidden.

(b) Boiling of Water

• When we heat water at 0°C (273 K), the particles have more energy compared to those in ice at the same temperature. As we continue to supply heat to the water, the particles move even faster. Eventually, they reach a point where they have enough energy to overcome the forces of attraction between them, and the liquid starts to change into a gas.
• The temperature at which a liquid begins to boil at atmospheric pressure is called its boiling point. Boiling is a process that occurs throughout the bulk of the liquid, where particles gain enough energy to change into the vapour state. For water, this temperature is 373 K (100°C). To convert temperatures between the Kelvin and Celsius scales, we subtract 273 from the Kelvin temperature to get the Celsius temperature, and vice versa. For example, 0°C = 273 K.
• The latent heat of vaporisation is the amount of heat energy required to change 1 kg of a liquid into gas at its boiling point and atmospheric pressure. This energy is absorbed by the particles in the form of latent heat, which is why particles in steam at 373 K (100°C) have more energy than those in water at the same temperature. 

(c) Sublimation
This shows that we can change the state of matter by altering the temperature. While most substances change from solid to liquid and then to gas when heated, some can change directly from solid to gas and vice versa without passing through the liquid state. This process is called sublimation when going from solid to gas, and deposition when going from gas to solid.

Lets understand Sublimation and Deposition of Camphor with an Activity 

• Take some camphor and crush it into small pieces. Place the crushed camphor in a china dish. Cover the dish with an inverted funnel.
• To the stem of the funnel, place a cotton plug. This setup will help in observing the process of sublimation and deposition.

Sublimation of camphor

• Observe the camphor over time. You will notice that the camphor starts to sublimate, which means it changes from solid to gas without passing through the liquid state. The gas will then deposit on the cooler parts of the funnel, changing back into solid. This process demonstrates sublimation and deposition.

2. Effect of Change of Pressure

By applying pressure, particles of matter can be brought close together

• Solid Carbon Dioxide (Dry Ice): Solid carbon dioxide, commonly known as dry ice, is stored under high pressure. When the pressure is reduced to 1 atmosphere, dry ice sublimates directly into gas without passing through the liquid state. This phenomenon occurs because of the specific conditions of pressure and temperature. 
• Influence of Pressure and Temperature: The state of a substance—whether it is a solid, liquid, or gas—is determined by the combination of pressure and temperature. By applying pressure and reducing temperature, gases can be liquefied. Conversely, changing the pressure and temperature can also lead to different states of matter. 
• Gaseous State under Pressure: When pressure is applied to a gas, the particles are forced closer together. This can lead to a change in the state of matter. For example, increasing pressure on a gas while reducing temperature can cause it to liquefy. 

Try yourself:

What can happen to matter?

• A.It can become invisible.
• B.It can disappear.
• C.It can grow.
• D.It can change its state.

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 a surface phenomenon because it occurs at the surface of a liquid only. Whatever the temperature at which evaporation takes place, the latent heat of vaporisation must be supplied whenever a liquid changes into a vapour (or gas).Evaporation

Explanation about Evaporation

• Some particles in a liquid always have more kinetic energy than 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 out to dry, 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:

1. 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.

2. 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.

3. Temperature: The Rate of evaporation increases with an increase in temperature. This is because with the increase in temperature number of particles gets enough kinetic energy to go into the vapour state (or gaseous state).
Example: Clothes dry faster in summer than in winter.

4. 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 higher, the rate of vaporisation decreases. The rate of evaporation is higher 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.

5. Wind speed: The rate of evaporation also increases with an increase in the speed of the wind. This is because with an increase in the speed of wind, the particles of water vapour move away with the 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 palms or surroundings and leave the palm feeling cool.
  (ii) We sprinkle water on the root or the open ground after a sunny, hot day. This cools the roof or open ground. This is because the large latent heat of vaporisation 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 bodies 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 the water within the pot falls, and hence it becomes cool.

Earthen Pot

• Rapid cooling of hot tea: If the 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 a 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.

Try yourself:

What process involves a liquid turning into a gas?

• A.Condensation
• B.Precipitation
• C.Sublimation
• D.Evaporation