1. Matter and its Classification

Chapter 1

Matter in our surroundings

Introduction

  • Everything in this universe is made of materials which scientist has names ‘matter’.
  • The matter is made up of very small tiny particles. It is not continuous but is particulate.
  • The matter is anything that occupies space and has mass.
  • Particles of matter have space between them and are continuously moving.
  • Particles of matter attract each other.

PHYSICAL NATURE OF MATTER

  1. Matter is made up of particles

If we take a 100 mL beaker, filling half of it with water and dissolve some salt/sugar. We will observe that there is no rise in water level and the salt/sugar has spread throughout the water (shown in the fig).

  1. How small are these particles of matter?
  • The particles of matter are very small – they are small beyond our imagination.

If we dissolve few crystals of potassium permanganate in about 1000mL of water, we will see the colour has changed. It shows that there must be millions of tiny particles in just one crystal of potassium permanganate, which keep on dividing themselves into smaller and smaller particles.

 

2. Molecular Theory of Matter

CHARACTERISTICS OF PARTICLES OF MATTER

1. Particles of matter have space between them

2.Particles of matter are co ntinuously moving –

  • They possess kinetic energy.
  • Increase in temperature also increases the kinetic energy of the particles
  • Thus, Particles of matter intermix on their own with each other.
  • They do so by getting into the spaces between the particles.
  • This intermixing of particles of two different types of matter on their own is called diffusion
  • On heating, diffusion becomes faster.

3. Particles of matter attract each other.

  • The strength of this force of attraction varies from one kind of matter to another.

3. Plasma and BEC State, Temperature

STATES OF MATTER

Matter around us exists in three different states– solid, liquid and gas, dependent on the characteristics of the particles of matter.

1. SOLID STATE – Solids have a definite shape, distinct boundaries and fixed volumes, that is, have negligible compressibility.

  • Solids may break under force but it is difficult to change their shape, so they are rigid.

          Examples, a pen, a book, a needle and a piece of wooden stick, a granule of sugar.

  • There are objects that are solid in state but seems do not follow the above rule but actually they do
  • A rubber band changes shape under force and regains the same shape when the force is removed. If excessive force is applied, it breaks.
  • A sponge has minute holes, in which air is trapped, when we press it, the air is expelled out and we are able to compress it.

2. LIQUID STATE – liquids have no fixed shape but have a fixed volume.

  • They take up the shape of the container in which they are kept.
  • Liquids flow and change shape, so they are not rigid but can be called fluid.
  • Solids, liquids and gases can diffuse into liquids (e.g. oxygen and carbon dioxide dissolves in water, which helps the survival of aquatic animals and plants)
  • The rate of diffusion of liquids is higher than that of solids.
  • This is because in the liquid state, particles move freely and have greater space between each other as compared to particles in the solid state.

3. GASEOUS STATE – gases are highly compressible as compared to solids and liquids

  • Due to its high compressibility, large volumes of a gas can be compressed into a small cylinder and transported easily
  • Examples: liquefied petroleum gas (LPG) cylinder, Compressed natural gas (CNG) fuel tanks
  • Due to high speed of particles and large space between them, gases show the property of diffusing very fast into other gases.
  • Rate of diffusion is much faster than solids and liquids
  • In the gaseous state, the particles move about randomly at high speed.
  • Due to this random movement, they exert pressure which is the force exerted by each gas particles per unit area on the walls of the container.

4. Interconversion of States of Matter

CAN MATTER CHANGE ITS STATE?

Yes, they do change under various circumstances and factors which we will study below

1. Effect of change of temperature

  • On increasing the temperature of solids, the kinetic energy of the particles increases.
  • Due to the increase in kinetic energy, the particles start vibrating with greater speed.
  • The energy supplied by heat overcomes the forces of attraction between the particles.
  • The particles leave their fixed positions and start moving more freely.
  • A stage is reached when the solid melts and is converted to a liquid.
  • The minimum temperature at which a solid melts to become a liquid at the atmospheric pressure is called its melting point.
  • The melting point of ice is 0° C (273.15 K).
  • The process of melting, is also known as fusion.
  • When a solid melts, its temperature remains the same
  • This increase in temperature (heat) of solids is used up in changing the state by overcoming the forces of attraction between the particles without showing any rise in temperature
  • This hidden heat is called latent heat
  • Therefore, the amount of heat energy that is required to change 1 kg of a solid into liquid at atmospheric pressure at its melting point is known as the latent heat of fusion.
  • So, particles in water at 0° C (273 K) have more energy as compared to particles in ice at the same temperature.
  • Similarly, when we supply heat energy to water, particles start moving faster.
  • At a certain temperature, a point is reached when the particles have enough energy to break free from the forces of attraction of each other.
  • At this temperature the liquid starts changing into gas.
  • The temperature at which a liquid starts boiling at the atmospheric pressure is known as its boiling point.
  • Boiling is a bulk phenomenon i.e. each particles of the liquid gain enough energy to change into the vapour state.
  • For water this temperature is 373 K (100°C = 273 + 100 = 373 K).
  • The input energy required to change the state from liquid to vapor at constant temperature is called the latent heat of vaporization
  • Water vapour at 373 K (100° C) have more energy than water at the same temperature. This is because particles in steam have absorbed extra energy in the form of latent heat of vaporization.

  • A change of state directly from solid to gas without changing into liquid state is called sublimation. Example, vaporization of camphor
  • The direct change of gas to solid without changing into liquid is called deposition. Example, soot in the chimney, making dry ice (solid carbon dioxide).

2. Effect of change of pressure

  • Applying pressure and reducing temperature can liquefy gases. Examples, LPG, liquid nitrogen.

5. Evaporation

EVAPORATION

  • In liquids, a small fraction of particles at the surface, having higher kinetic energy, is able to break away from the forces of attraction of other particles and gets converted into vapour. This phenomenon of change of a liquid into vapours at any temperature below its boiling point is called evaporation.

1.Factors affecting evaporation

  • The rate of evaporation increases with
  • increase of surface area
  • increase of temperature
  • decrease in humidity
  • increase in wind speed

1. How does evaporation cause cooling?

The particles of liquid absorb energy from the surrounding to regain the energy lost during evaporation. This absorption of energy from the surroundings makes the surroundings cold.

-Why should we wear cotton clothes in summer?

* During summer, we perspire more because of the mechanism of our body which keeps us cool. We know that during evaporation, the particles at the surface of the liquid gain energy from the surroundings or body surface and change into vapour. The heat energy equal to the latent heat of vaporisation is absorbed from the body leaving the body cool. Cotton, being a good absorber of water helps in absorbing the sweat and exposing it to the atmosphere for easy evaporation.

-Why do we see water droplets on the outer surface of a glass containing ice-cold water?

* If we take some ice-cold water in a tumbler. Soon water droplets forms on the outer surface of the tumbler. The water vapour present in air, on coming in contact with the cold glass of water, loses energy and gets converted to liquid state, which we see as water droplets.

 

1. Pure Substance

Is Matter around us pure

 

Introduction

When we talk about pure, it means that all the constituent particles of that substance are the same in their chemical nature. A pure substance consists of a single type of a particles. What is the type of pure substances?

  • Elements

 

  • Robert Boyle A was the first scientist to use the term element in 1661.
  • Antoine Laurent Lavoisier (1743–94), a French chemist, was the first to establish an experimentally useful definition of an element.
  • Elements can be normally divided into metals, non – metals and metalloids.

Metals

  • Metals usually show some or all of the following properties:
  • They have a lustre (shine).
  • They conduct heat and electricity.
  • They are ductile (can be drawn into wires). 
  • They are malleable (can be hammered into thin sheets).
  • They are sonorous (make a ringing sound when hit). 

# Examples of metals are gold, silver, copper, iron, sodium, potassium etc. 

# Mercury is the only metal that is liquid at room temperature.

Non metals

  • Non – metals usually show some or all of the following properties:
  • They are poor conductors of heat and electricity.
  • They are not lustrous, sonorous or malleable. 

# Examples of non – metals are hydrogen.  oxygen, iodine, carbon (coal, coke).  bromine, chlorine etc.

Metalloids

  • Metallaoids have intermediate properties between of metals and non – metals. 

# Examples are boron, silicon, germanium etc.

Mixture and compound

Mixture Compound
1. Elements or compounds are simple calling so new substance is formed. Compound   1. Substances Are Reated Together with each other to make a new substance.  
2. Elements do not combine in a fixed ratio.   2. Compositions the the component is Fixed i.e. , They combine together in a fixed ratio according to their masses.  
3. A mixture shows the properties of its components   3. compound does not show the Properties of component elements.  
4. Components can be easily separated by any mechanical method which is suitable.   4. components can not be separated from each other by simple mechanical methods.

 

2. Impure Substance

But when we see around us, we observe most of the matter around us exists as mixtures of two or more pure components. For example: Sea water, Air etc.

WHAT IS A MIXTURE?

 It is a form of matter in which two or more elements or compounds combine physically in any proportion by weight.

Characteristics of Mixture

  • Mixture may be homogeneous and heterogeneous.
  • Mixture does not have a fixed melting point.
  • In a mixture, the different constituents combine physically in any proportion by mass.
  • The constituents of a mixture do not lose their identical property.
  • Usually, no energy change take place during the formation of a mixture.

Types of Mixtures

  • Homogeneous mixture: A mixture which has same composition throughout. It has no visible boundaries of separation between the various constituents Solutions are homogeneous mixtures. For example, Detergent in water, Sugar in water, Ice cream etc.
  • Heterogeneous mixture: A mixture which has different compositions in different parts. These types of mixtures have visible boundaries of separation between the various constituents. For example, Oil in water, Fruit salad, Sand in water etc.

3. Separation of Mixtures

WHAT IS SOLUTION?

A solution is a homogenous mixture of two or more substances. E.g.  Nimboo pani, Soda water. A solution has a solvent and a solute as its components. 
Solvent: The component of the solution that dissolves the other component in it is called the solvent.
Solute: The component of the solution that is dissolved is called the solute. in the solvent
Alloys: They are the mixtures of two or more metal or a metal and a non-metal and cannot be separated into their components by physical methods.

Properties of a Solution

•    A Solution is a Homogeneous mixture.
•    The particles of a solution are smaller than 1nm (10^-9 Metre) in Diameter. They cannot bean by Naked eyes.
•    They are of very Small Particle Size, so they do not scatter a beam of lighting passing through the solution.
•    The Solute Particles Cannot be separated from the mixture by the process of filtration.

Concentration of a solution

•    Saturated solution: Depending upon the amount of solute present in a solution, it can be called a dilute, concentrated or a saturated solution. 
•    Unsaturated solution: If the amount of solute contained in a solution is less than the saturation level, it is called an unsaturated solution.
•    Solubility: The amount of the solute present in the saturated solution at this temperature is called its solubility. 
•    The concentration of a solution is the amount of solute present in a given amount (mass or volume) of solution, or the amount of solute dissolved in a given mass or volume of solvent.
•    Concentration of solution = Amount of solute / Amount of solution
                                            Or

Amount of solute / Amount of solvent

Ways of expressing the concentration of a solution

 

•    Mass by mass percentage of a solution

Mass of solute / Mass of solution x100
•    Mass by volume percentage of a solution

Mass of solute / Volume of solution x100
 

 

4. Physical And Chemical Changes

What is a suspension?

In which solids are dispersed in liquids, are called suspensions. 
A suspension is a heterogeneous mixture
Particles of a suspension are visible to the naked eye.
Properties of a Suspension

      •    Suspension is a heterogeneous mixture.
      •    The particles of a suspension can be seen by the naked eye.

  • The particles of a suspension scatter a beam of light passing through it and make its path visible.

      •    The solute particles settle down when a suspension is left undisturbed, that is, a suspension is unstable.

WHAT IS A COLLOIDAL SOLUTION?

    A colloidal solution is a heterogeneous mixture, for example, milk. 
    Because of the small size of colloidal particles, we cannot see them with naked eyes.
    These particles can easily scatter a beam of visible light.
        Tyndall effect
    The scattering of a beam of light is called the Tyndall effect
    The Tyndall effect can also be observed when a fine beam of light enters a room through a small hole. 
    This happens due to the scattering of light by the particles of dust and smoke in the air.

Observation of Tyndall effect

    The Tyndall effect can be observed when sunlight passes through the canopy of a dense forest. 
Properties of a colloid.
    A colloid is a heterogeneous mixture. 
    The size of particles of a colloid is too small to be individually seen by naked eyes.
    Colloids are large enough to scatter a beam of light passing through it and make its path visible.
    They do not settle down when left undisturbed, that is, a colloid is quite stable.
    They cannot be separated from the mixture by the process of filtration. 


Dispersing medium

    The components of a colloidal solution are the dispersed phase and the dispersion medium.
    The solute – like component or the dispersed particles in a colloid form the dispersed phase, and the component in which the                dispersed phase is suspended is known as the dispersing medium.


 

1. Dalton's atomic Theory

Chapter 3

Atoms & Molecules

 

Dalton’s Atomic Theory

The matter is made up of indivisible particles known as atoms.

The properties of all the atoms of a given element are the same, including mass. This can also be stated as all the atoms of an element have identical mass and chemical properties; atoms of different elements have different masses and chemical properties.

Atoms of different elements combine in fixed ratios to form compounds.

Atoms are neither created nor destroyed. The formation of new products (compounds) results from the rearrangement of existing atoms (reactants) in a chemical reaction.

The relative number and kinds of atoms are constant in a given compound.

 

 

Laws of Chemical Combination

Given by Lavoisier and Joseph L. Proust as follows:

 

Law of conservation of mass

According to the law of conservation of mass, matter can neither be created nor destroyed in a chemical reaction. It remains conserved.

Mass of reactants will be equal to the mass of products.

 

Law of constant proportions

A pure chemical compound contains the same elements combined together in a fixed proportion by mass is given by the law of definite proportions.

For e.g., If we take water from a river or from an ocean, both have oxygen and hydrogen in the same proportion.

 

 

Atom
Atoms are the smallest particles of an element which can take part in a chemical reaction.
Size of an atom: atomic radius is measured in nanometres.

The atomic symbol has three parts: -

The symbol X: the usual element symbol

The atomic number A: equal to the number of protons

The mass number Z: equal to the total number of protons and neutrons in an element.

 

2. IUPAC and atomic Symbols

                                              Atom
Atoms are the smallest particles of an element which can take part in a chemical reaction.
Size of an atom: atomic radius is measured in nanometres.
The atomic symbol has three parts: -
The symbol X: the usual element symbol
The atomic number A: equal to the number of protons
The mass number Z: equal to the total number of protons and neutrons in an element.
                                                      Atomic Mass

Atomic mass and atomic mass unit

Atomic mass is the total of the masses of the electrons, neutrons, and protons in an atom, or in a group of atoms, the average mass.
Mass of an atomic particle is called the atomic mass.
This is commonly expressed as per the international agreement in terms of a unified atomic mass unit (AMU).
It can be best defined as 1/12 of the mass of a carbon -12 atom in its ground state.
Molecule

It is the smallest particle of an element or a compound which can exist independently.
•    Molecules of an element constitute the same type of atoms.
•    Molecules may be monoatomic, diatomic or polyatomic.
•    Molecules of compounds join together in definite proportions and constitute a different type of atoms.
 

 

 

 

 

 

3. Compound

                                               Atomicity
The number of atoms constituting a Molecule is known as its atomicity.

Compounds

A pure substance made up of two or more elements chemically combine together in a fixed ratio under fixed condition is called compound. Example: Calcium carbonate, Common salt, Sugar. 

Properties of compound

 1.Compounds can be separated into the constituent only by chemical methods.  
2. Properties of compound differ from the properties of their constituents. 
3. During the formation of a compound energy is absorbed or released.
4. The compound are homogeneous.

 

Valency

The combining capacity of an element is known as its valency. Valency is used to find out how the atom of an element will combine with the atom of another element to form a chemical compound.
(Every atom wants to become stable, to do so it may lose, gain or share electrons.)

•    If an atom consists of 1, 2 or 3 electrons in its valence shell then its valency is 1, 2 or 3 respectively,
•    If an atom consists of 5, 6 or 7 electrons in the outermost shell, then it will gain 3, 2 or 1 electron               respectively and its valency will be 3, 2 or 1 respectively.
•    If an atom has 4 electrons in the outermost shell than it will share this electron and hence its valency         will be 4.
•    If an atom has 8 electrons in the outermost electron and hence its valency will be 0.

Ions
An ion is an electrically charged atom or group of atoms.
An ion is formed by the loss or gain of electrons by an atom, so it contains an unequal number of electrons and protons.

There are two types of ions:

Cation: A positively charged ion is known as cation.
For Ex: Na+, Mg2+ 
A cation is formed by the loss of one or more electrons by an atom.

Na – e– ——–>Na+
Z=11                 Z=10
2,8,1                  2,8
K, L, M                  K, L
 

Mg – e– ——–> Mg2+
Z=12                 Z=10
2,8,2                  2,8
K,L,M                  K,L
 

Anion: A negatively charged ion is known as anion.
For Ex: Cl–, O2- 
An anion is formed by gain of one or more electron by an atom.

Cl + e– ——–> Cl–
Z=17                 Z=18
2,8,7                  2,8,8
K, L, M                 K, L, M
 

O + e– ———> 02-
Z=8                   Z=10
2,6                    2,8
K, L                    K, L
 

Radicals

An atom or group of atoms having a charge, i. e. either negative or positive, on it.
The radicals having positive charge are called cations.eg. Sodium ion.
The radicals having negative charge are called anions. eg. Chloride ion.








 

4. Chemical Formula

                                   Chemical Formulae

Rules:

(i) The valencies or charges on the ion must balance.

(ii) Metal and non-metal compound should show the name or symbol of the
metal first.
e.g., Na+ Cl– → NaCl

(ii) If a compound consists of polyatomic ions. The ion is enclosed in a bracket before writing the number to indicate the ratio.
e.g., [SO4]2- → polyatomic radical
H1+ SO42- → H2SO4

Molecular Mass
It is the sum of the atomic masses of all the atoms in a molecule of the substance. It is expressed in atomic mass unit (u).

Formula Unit Mass
It is the sum of the atomic masses of all atoms in a formula unit of a compound. The constituent particles are ions.


 

 

Mole Concept

Definition of mole: It is defined as one mole of any species (atoms, molecules, ions or particles) is that quantity in number having a mass equal to its atomic or molecular mass in grams.
1 mole = 6.022 x 1023 in number
Molar mass = mass of 1 mole → is always expressed in grams and is also known as gram atomic mass.
l u of hydrogen has → 1 atom of hydrogen and 1g of hydrogen has → 1 mole of hydrogen
= 6.022 x 1023 atoms of hydrogen.
 

1. Introduction

Chapter 4

Structure of an Atom

Charged particles in Matter

•    Atoms are the basic building blocks of matter.
•    Different kinds of matter exist because there are different kinds of atom
 present in them.

Postulates of Dalton’s Atomic Theory

•    All matter is made up of tiny, indivisible particles called atoms.
•    All atoms of a specific element are identical in mass, size, and other properties. However, atoms of different element exhibit different properties and vary in mass and size.
•    Atoms can neither be created nor destroyed. Furthermore, atoms cannot be divided into smaller particles.
•    Atoms of different elements can combine with each other in fixed whole-number ratios in order to form compounds.
•    Atoms can be rearranged, combined, or separated in chemical reactions.
•    John Dalton considered atom to be an individual entity, but his concept had to be discarded                  at the end of 19th century, when scientist through experiments able to find existence of charge (electrons & protons) and neutral particles (neutrons) in the atom. These particles were called sub atomic particles.
•    The discovery of electron and proton is credited to J.J. Thomson and E. Goldstein, respectively. 
•    J.J. Thomson proposed that electrons are embedded in a positive sphere
•    Electron was represented as 'e-' and proton as 'p+'. The mass of a proton is taken as one unit and its charge as plus one where the mass of an electron was considered to be negligible and its charge minus one
•    It seemed that atom consisted of electrons and protons which balanced their charges mutually.

 

2. Discovery of Electrons and Protons

                                   The Structure of an atom

• Dalton's atomic theory suggested that the atom was indivisible and indestructible. However, the discovery of two fundamental particles in the atom the electrons and protons led to the failure of this aspect of the theory.
 • Hence, J.J. Thomson was the first to propose a model for the structure of an atom.

Thomson’s Model of an Atom

According to J.J. Thomson, the structure of an atom can be compared to Christmas pudding where electrons are present inside a positive sphere.


 

 

 

 

 

 

•    An atom is composed of a positively charged sphere in which electrons are embedded.
•    Atom is neutral as the positive and negative charged are equal in proportion.
Rutherford’s Model of an Atom

Rutherford’s Experiment

•    He experimented with thin gold foil by passing alpha rays through it.
•    He expected that the gold atoms will deflect the Alpha particles.

 

 

 

 

 

 

 

Observations

Inferences

Alpha particles which had high speed moved straight through the gold foil

Atom contains a lot of empty space

Some particles got diverted a by slide angles

Positive charges in the atom are not occupying much of its space

Only one out of 12000 particles bounced back

The positive charges are concentrated over a particular area of the atom.

Thus, Rutherford gave the nuclear model of an atom based on his experiment which suggests that -

•    Atoms contain a lot of unoccupied space
•    There is a heavily positively charged substance present in the center of the atom which is called the          nucleus
•    The nucleus contains an equal amount of positive and negative charge.


 

 

 

 

 

 

 

Drawbacks of the Nuclear Model of an Atom

•    The Nuclear Model of the Atom failed to explain how an atom remains stable despite having positive and negative charges present in it. 
•    Maxwell has suggested a theory according to which if any charged particle moves in a circular motion, it radiates energy.
•    So, if electrons start moving in a circular motion around the nucleus, they would also radiate some energy which would decrease at the speed of the electrons. 
•    As a result, they would fall into the nucleus because of its high positive charge.

What are nucleons? –  Protons and Neutrons are collectively called as Nucleons.

3. Discovery of Neutrons

Bohr Model of an atom

Bohr suggested that –

•    Electrons spin around the nucleus in an individualized separate path or unattached          orbit.
•    The electrons do not emit any energy while moving Indies special orbits.
•    These orbits are also called as Energy Levels.
•    They are represented using letters or numbers as shown in the figure below –


 

 

 

 

 

 

 

The Neutrons

J. Chadwick discovered that there is another sub-atomic particle present in the atom. This particle carries no charge and is known as a Neutron. Therefore, we can conclude that atom consists of three types of particles –

Electrons

which carry a negative charge

Protons

which carry a positive charge

Neutrons

they are neutral 

 

The distribution of electrons in different shells or orbits

•    If Orbit number = n
•    Then number of electrons present in an Orbit = 2n2
•    So, for n =1
•    Maximum electrons present in shell – K = 2 * (1)2 = 2
•    The outermost shell can contain at most 8 electrons.
•    The shells in an atom are filled in sequence.
•    Thus, until the inner shells of an atom are filled completely the outer shells cannot         contain any electrons.

 

4. Some Important Defintion

Valency

  •    Valence Electrons – Electrons existing in the outermost orbit of an atom are called Valence Electrons.

•    The atoms which have completely filled the outermost shell are not very active chemically.
         •    The valency of an atom or the combining capacity of an atom is given by the number of elements present in the outermost shell.

        •    For Example, Helium contains two electrons in its outermost shell which means its valency is two. In other words, it can share two electrons to form a chemical bond with another element.

           •    What happens when the outermost shell contains a number of electrons that are close to its maximum capacity?

Valency in such cases is generated by subtracting the number of electrons present in the outermost orbit from octet (8). For example, oxygen contains 6 electrons in its outermost shell. Its valency is calculated as: 8 – 6 = 2. This means oxygen needs two electrons to form a bond with another element.

Atomic Number of an Element

Atomic Number (Z) = Number of protons in an atom
Mass Number of an Element
Mass Number = Number of protons + Number of neutrons


 


 

 

 

 

 

 

Isotopes

•    The atoms of an element can exist in several forms having similar atomic numbers but varying mass numbers.
•    Isotopes are pure substances.
•    Isotopes have a similar chemical nature.
•    Isotopes have distinct physical characteristics.

 

 

 

 

 

 

 

Where can we use Isotopes?

1. The fuel of Nuclear Reactor – Isotope of Uranium
2. Treatment of Cancer – Isotope of Cobalt
3. Treatment of Goitre – Isotope of Iodine

Isobars

The atoms of several elements can have a similar mass number but distinct atomic masses. Such elements are called Isobars. 

Isotones

•    Species having same number of neutrons but different number of protons are called Isotones.


 

 

 

 

•    Examples include boron-12 and carbon-13 nuclei both contain 7 neutrons, and so are isotones. 

Difference Between Isotopes, Isobars & Isotones

 

 

Related Unit Name