Nature of matter

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					Unit – Matter: Nature and Behaviour (45 periods) Marks allotted- 18

Chapters

    1) Nature of matter: Classification of matter based on chemical constitution- elements,
       compound and mixtures, types of mixtures- homogeneous and heterogeneous solution,
       suspension and colloid, concentration of solution (percentage only)
                Atoms and molecules, atomic theory of matter(Dalton’s postulates), atomic and
       molecular masses, the mole, Law of constant proportion, calculation of % composition of
       elements in simple compounds, determination of empirical and molecular formulae of
       simple substances.



                                    Nature of matter

Introduction

Chemistry is based on how matter reacts and combines. So in this Chapter we will study about
Matter and related things. This chapter is very important from examination point of view. This
chapter includes theory as well as numericals. It provides you information about matter and its
constituents. It shows us the similarity and differences between different objects. It will give you
some interesting facts from our day-to-day life.

        Why it’s not possible to compress a solid?
        How many states of matter are there?
        What is Plasma?
        Which metal is liquid at room temperature?
        Which non-metal is liquid at room temperature?
        What are metalloids?
        What is an atom?
        Why you won’t be able to dissolve sugar in water after some time?
        When you pour some water in a container, why liquid always fill the bottom of the
         container?
        When you seat under a fan after sweating, why you feel so relax and cool?
        What is mist?
        Why you feel the smell of perfume at every corner of the room, when a bottle of perfume
         is opened in a room?
        What is common salt? Is it a compound or mixture?
        Why water is considered as a compound not as a mixture?
        Why milk is considered as a heterogeneous mixture?
        Why it is considered that medicines are more effective in colloidal state?
        What is the role of colloids in delta formation?




For finding out the answers to above questions, carefully go through the chapter.


Matter
In our daily life we see large variety of objects around us. The nature and behaviour of different
types of substances around us are different. These substances differ in their appearance, colour,
taste, smell, texture etc. For example salt. Sugar. Water, oxygen, carbon dioxide etc.

Do you ever thought that all things around us are alike in one way. They all have mass and
occupy space that’s why they all are matter. In fact, everything in this world consists of matter and



Matter is anything that occupies space and has mass.

Thus, matter could be any thing — living or non-living around us, from a chair to a tree, from a
stone piece to a mountain, from an almirah to a glass bottle, from a concrete pillar to a building.
Therefore, we can say that all the objects around us are composed of matter.

 In ancient India, Kanada (about 600 BC) was the first philosopher who stated that everything in
this world is made of the five basic elements: earth, fire, air, water and sky. These elements could
undergo any transformation into various substances depending upon their main qualities and
composition. Kanada was the first to propose that all the matters of this universe are made up of
eternal and imperceptible particle known as paramanu.

Thus, to understand the world around you, it is essential to know about matter.

Classification of Matter
Matter can be classified in two principal ways:

    1. On the basis of its physical state as solid, liquid and gas
    2. On the basis of its chemical constitution as an element, compound and mixture.


Physical states of matter
    Matter can be classified into following three categories according to its physical nature: Solid,
    Liquid and Gas.

            a.       Solid: Solid is a physical state of matter that has definite shape and volume.
                            Your book, Pen, Pencil, Chalk, Board all are solids.

            Properties of the solids:

                          Solids are rigid and have a definite shape.
                          Solids have definite size, shape and the volume and it is independent
                           of shape, size and volume of the container.
                          Solids are nearly incompressible (A property of a matter because of
                           which matter resists to the change of shape, size and volume due to
                           pressure).
                          Solids diffuse very slowly in comparison to liquids and gases.
                          Most of the solids are crystalline in nature.
                          Solids have high densities. They are heavy.
                          Solids do not fill the container completely.



   Have you ever tried to compress a solid? What was the result of that compression?

   You have noticed that solids are nearly incompressible because when you compress
   something you are pushing the atoms of the substance closer to each other. When pressure
   continues substances are compressed. Solids already have their atoms close together so
   it's hard to push them even closer.
           b.           Liquid: Liquid is a physical state of matter that has no definite shape but has
                        definite volume. The best example, which we can provide here, is water.




                Solid        +      Heat energy                    Liquid



When you pour some water in a cup, it fills the bottom of the container first and then fills the
rest. Do you know why liquid always fill the bottom of the container and why it has a flat
surface in a container?

Liquid starts filling at the bottom because of gravity. When it is in that container it also has a
flat surface. That’s because of gravity too.



           Properties of the liquids:

                                Liquids have fluidity and do not have a definite shape.
                                The shape of the liquid depends upon the shape and size of the
                                 container in which they exist.
                                Liquids are slightly compressible.
                                Liquids diffuse faster than solids but slower than gases.
                                Liquids have moderate to high densities. They are usually less dense
                                 than solids.
                                Liquids do not fill the container completely.




           c.           Gas: Gas is a physical state of matter that has no definite shape or volume.
                        For example, if you wave your arms around you can find a gas (air).

           Properties of the gas:

                                Gases are more fluid and do not have definite shape or size.
                                The shape of the gas depends upon the shape and size of the
                                 container in which they exist.
                                Gases are more compressible.
                                Gases diffuse faster than both liquids and solids.
                                Gases have very low densities. They are very, very light.
                                Gases fill their container completely.
These three physical states of matter can be explained on the basis of different types of
intermolecular force between constituent particles (atom or molecule) that hold them together.

In solids, these forces are very strong and thus, keep the particles closer to each other, which is
the reason why the solids have definite shape, high density and incompressibility (it means the
solids can sustain against the high pressure and it affects little on the volume). The ordered
arrangement of the constituent particles in the solids is known as lattice and this gives a definite
geometrical shape to the solids. Thus, the main identifying characteristic of solids is their rigidity.


In liquids, these forces are not strong enough to keep the constituent particles in the fixed
position. As a result of this, liquids have fluidity and small compressibility (it means high pressure
affects its volume slightly).


In gases, these forces are extremely weak and the constituent particles are free to move in about
any direction and can occupy any space available to them, thus they can occupy entire volume of
the container. The particles collide with each other and also with the walls of the container. Thus,
they exert a steady force on the surface of the wall and results into the pressure of gas. For
example, natural gas is compressed and is supplied as a fuel for vehicles in the name of CNG
(Compressed Natural Gas).




Do you Know, In 1879 Sir William Crookes, an English physicist, identified a fourth state of matter,
now called plasma. Plasma is by far the most common form of matter. Plasma makes up over
99% of the visible universe and perhaps most of that which is not visible. Neon in neon lights is
also in plasma state.




Basic characteristics of the three physical state of matter

                                     Solids                 Liquids                     Gases
Shape and Size            Definite                  Indefinite              Indefinite
                                                    (Takes the shape of     (Spread all over the
                                                    container in which they container)
                                                    are placed)
Density                   Higher than liquids       Higher than gases but      Lower than both solids
                          and solids                lower than solids          and liquids
Compressibility           Very less                 Slightly compressible      Compressible
                          compressible
                         (negligible)
Fluidity or Rigidity     Rigid                    Fluid                    Fluid
Intermolecular forces Very strong                 Stronger than gases     Weaker than both
                                                  but weaker than solids. solids and liquids.
Diffusion                Do not diffuse           Diffuse slowly           Diffuse rapidly

Many kinds of matter exist in all these three states under appropriate conditions of temperature
and pressure. For example, we know that ice (solid), water (liquid) and steam (gas) have the
same matter that exists in different physical state.




Chemical Change and Physical Change


Chemical Change:

A change in which one or more substances changes chemically by forming new substance.
Chemical changes include chemical reaction
For example
        C  O2          CO2
        2Na  Cl2        2NaCl

In the above chemical reactions, reactants are different from products in properties like carbon
and oxygen has their own individual properties and properties of carbon dioxide are entirely
different.

Physical Change:

 A change in which a substance is changing physically and no new compound substances is
formed. In physical change, the product formed is chemically alike with the starting substance.

For example: On freezing water (liquid) changes into ice (solid) and on melting ice will change to
water (liquid)

Similarly, in a sugar solution, sugar and water both have their independent properties (they retain
their properties in solution). Sugar can be recovered by evaporation.




Effect of Temperature and Pressure on the States of Matter (Change of
State):



A substance may exist in any of the three different states of matter (i.e., solid, liquid or gas)
depending upon the conditions of temperature and pressure. Here, we can take the example of
water, when ice is heated it changes into water and if we continue the process of heating water
results into steam.

Ice       Heat                   water         heat                   steam

(solid)                         (liquid)                            (gas)




Effect of temperature on solids:

When you heat any solid, it changes into liquid because when the solid is heated, it initially
expands as the intermolecular force between the particles weakens and the particles due to the
newly found kinetic energy start vibrating. As the heating progresses, the constituent particles
start vibrating more vigorously in their respective positions. When the supplied heat or thermal
energy is just sufficient to overcome the intermolecular force between the constituent particles,
the solid changes into the liquid. This process of changing over from solid to liquid is known as
melting. The temperature at which melting occurs under normal pressure is called the melting
point of the substance. Every solid has a characteristic melting point.




Effect of temperature on liquids:

 You must have been noticed that when you heat water you find some vapours. Actually when
you heat water or any liquid you are providing energy to their atoms and molecules that’s why
they get excited and escape from the liquid, Thus forming a gas. Under normal pressure
conditions, liquid on heating changes to gas. This process is known as evaporation. The
constituent particles on the surface of the liquid on evaporation exert pressure on the surface of
the liquid, as they are unable to push their way into the atmosphere because of their low pressure
inside the liquid in comparison to the pressure of atmosphere. This pressure exerted by the
constituent particles in the gaseous state on the surface of the liquid is known as vapour
pressure. As the heating progresses, the constituent particles gain more and more kinetic energy
and results in more particles entering into the evaporation phase and at one stage, the vapour
pressure over the surface becomes equal to the atmospheric pressure. All the constituent
particles enter into the gaseous state. The temperature at which this change occurs from liquid to
the gaseous state is called the boiling point of the liquid. Every liquid has a characteristic boiling
point.




      Do you know when you seat under a fan after sweating, why you feel so relax and cool?
      You feel cool and relax because of evaporation. Evaporation brings cooling in the
      environment.



Effect of temperature on gases:

Molecules of gases are full of energy and they continuously bounce. When you decrease the
temperature energy will be sucked out of gaseous atoms and molecules, they result in the
formation of liquid. This process of changing gas into liquid is known as Condensation. For
example, you must have notice that while cooking something in water, when you lift the lid of your
container, some water droplets are there on the lid. Theses water droplets are formed from
vapours /steam. If the temperature drop further to freezing point it leads to drop in energy in the
molecules and attraction forces allow the molecules to group together, resulting in the formation
of a solid.




 During winter you must have noticed mist (water droplets) on substances. Do you know from
 where this drops come from? Yes, your answer is correct if it is vapour. Actually due to low
 temperature water vapours present in air converts into water droplets. (Condensation)



Gases undergo diffusion or effusion. Diffusion is a process in which in any gaseous mixture kept
at a uniform temperature, the composition eventually becomes uniform throughout the system, no
matter what was the original solution. This is explained by the diffusion that is one form of
molecular motion. While effusion is a process in which the gas molecules spread out into the
atmosphere. For example, when a bottle of perfume is opened in a room, a person can feel the
smell of the perfume at any corner of the room.




                                                             Gases

                             Liquids

Solids                                       Energy

         Energy




On heating, the constituent particles of gas gain kinetic energy and move randomly. In the
process, they collide with each other and with the walls of container and thus exert pressure.

Effect of pressure on solids: Solids can be classified as true solids as they have rigid structure
and definite shape. When the solids are compressed, they do not distort and resist any change
due to compression.

Effect of pressure on liquids: When pressure is applied to liquids, they get compressed slightly.
This is because liquids have a very little free space between their constituent particles.
Effect of pressure on gases: Any gas kept in a closed container exerts pressure on the walls of
the container. As the free space between the constituent particles is more. On applying pressure,
gases get compressed easily. On increasing the pressure, they exert so much pressure that the
walls of the container might burst as we can see in the case of balloon.




Chemical Constitution of Matter

Matter can be classified into following three categories according to its chemical constituent:
Element, Compound and Mixture

    i.       Element: Element is a substance, which cannot be decomposed into simpler substances
             by ordinary chemical methods or simple physical process.

             An element is made up of atoms, all having the same atomic number. For example,
             hydrogen is an element, as it cannot be split up into two or more simpler substances by
             the ordinary methods like applying heat, light or electricity. All the atom of hydrogen has
             the same atomic number 1. Similarly, oxygen is an element because it cannot be broken
             down into simpler substances and all the atoms of oxygen have the same atomic number
             8. Water is made up of 2 hydrogen and1 oxygen and it can be split into two simpler
             substances hydrogen and oxygen by using electricity (electrolysis), therefore its not an
             element.


                  There are 115 elements are known at present, out of which 92 elements occur in
                  nature, while the remaining 23 elements have been prepared artificially. Some of
                  the common elements are: hydrogen, oxygen, nitrogen, carbon, sulphur, chlorine,
                  bromine, gold, silver, Zinc, iron and mercury.




Element can be classified into following three categories according to its physical nature: Solid,
Liquid and Gas.

         a. Solids like carbon, gold, silver etc.
         b. Liquids like bromine and mercury
         c. Gases like hydrogen, nitrogen, oxygen etc.

         On the basis of properties elements can be of three types:
         a. Metals like Magnesium, iron and calcium.

              All metals are solid except mercury, which is liquid.

         b. Non metals like hydrogen, carbon and oxygen

          Bromine is the only liquid non-metal.

     c. Metalloids: These elements have characteristics common to metals as well as non-
     metals. Tin, Bismuth, antimony are common examples.


.         80% of elements are metals and remaining 20% are the non-metals
Symbols of elements

In order to simplify, element are represented by symbols. The symbol of an element is the first
letter and another letter of the English name or latin name of the element. For example,
hydrogen is represented as H, carbon as C, iron as Fe, copper as Cu, etc.



Symbols Derived from English names of the Elements

Element                     Symbol                   Element                  Symbol
Hydrogen                    H                        Helium                   He
Lithium                     Li                       Carbon                   C
Nitrogen                    N                        Oxygen                   O
Flourine                    F                        Neon                     Ne
Neon                        Ne                       Magnesium                Mg
Aluminium                   Al                       Silicon                  Si
Phosphorous                 P                        Sulphur                  S
Chlorine                    Cl                       Argon                    Ar
Calcium                     Ca                       Manganese                Mn
Nickel                      Ni                       Zinc                     Zn
Bromine                     Br                       Iodine                   I


Symbols Derived from Latin names of the Elements

Element                                              Symbol
Sodium                                               Na
Potassium                                            K
Iron                                                 Fe
Copper                                               Cu
Silver                                               Ag
Gold                                                 Au
Mercury                                              Hg




  ii.       Compound: Compound is a substance, which is produced by the chemical reaction of
            two or more elements. A compound is a pure substance made up of two or more
            elements chemically combined with one another in a fixed proportion by mass. For
            example, water consists of 2 hydrogen and 1 oxygen. Other examples are: carbon
            dioxide, sodium chloride, etc.




        2 atoms of Hydrogen + 1 atom of Oxygen                      1 molecule of Water

        Most of the materials in nature are mixtures of elements and compounds in different
        proportions.
 iii.      Mixture: Mixture is a substance, which is formed by mixing two or more substances
           which chemically do not react with each other, but can be mixed in any proportion. For
           example, air is a mixture of several gases. Other examples of mixtures are: seawater,
           milk, mortar, etc.


        Mixtures can be homogeneous or heterogeneous.

                  A homogeneous mixture has uniform composition throughout its mass. The
                   various components cannot be visually distinguished. For example, air, sugar
                   solution, alloys like steel (iron + carbon), brass (copper + zinc), bronze (copper
                   and tin).

               Water +      Sugar                 Solution of sugar in water (Homogeneous Mixture)




                  A heterogeneous mixture does not have a uniform composition throughout its
                   mass. The various constituents can easily be distinguished. For example, mixture
                   of sand in water, mixture of sand in sugar, suspension of mud in water. Almost all
                   the mixtures are heterogeneous except for solutions and alloys.




               Water   +    Sand                  Solution of sand in water (Heterogeneous Mixture)




 Do you know Milk is also a heterogeneous mixture? But when you see milk, it appears
 to be homogeneous mixture. However, when we observe it with help of microscope, it
 appears as drops of fat are suspended in clear liquid. Thus, milk is also a
 heterogeneous mixture.



It is easier to study compounds and mixtures by comparing their properties.

                       Compounds                                            Mixtures
Properties of compounds are different from those of       Individual properties of the constituents are
its constituents. For example, Sodium is a metal,         retained in a mixture. For
while sodium chloride is a salt.                          example, Iron fillings and sand.
Elements in a compound are in a fixed proportion by Elements in a mixture can be mixed in any
weight.                                             proportion.
                                                          Physical processes (such as evaporation,
Compounds cannot be split-up into their constituents      sublimation, distillation, magnet, solvents,
by simple processes. For example, in sodium               etc.) can separate the constituents of a
chloride, sodium and chlorine can be separated only       mixture.
by chemical processes.                                    For example, mixture of iron fillings and
                                                          sand can be easily separated by a magnet.
Energy is either released or absorbed when a               No energy transfer is involved in the
compound is formed or decomposed.                          formation (or separation) of a mixture.
                                                           Their melting point, boiling points, etc.
Their melting point, boiling point, etc. do not change.
                                                           keeps on changing.
                                                           Mixtures may be homogeneous or
A compound is homogeneous substance.
                                                           heterogeneous.




                                       Matter

                                Solid, Liquid or Gas




    Heterogeneous Matter                                  Homogeneous Matter
    (Variable Composition)                                (Uniform Composition)
    Soil, Air                                             Salt. Water, hydrogen




                                                  Solutions                              Pure substances
                                          (Non-fixed composition)                    (Fixed Composition)
                                          Homogeneous Mixtures                       Water, sugar, hydrogen
                                          Sugar in water




                                                                        Elements
                                                                        (Cannot be                     Compounds
                                                                        decomposed)                    (Can be
                                                                                                       decomposed)
Solution

Solution is a homogeneous mixture of two or more substances, which are chemically non-
reacting.
In our day-to-day life we see various forms of solution. It may be noted that that most of the
common solutions contain solid solutes and liquid solvents. For example, sugar is dissolved in
water to prepare sharbat and iodine is dissolved in ethyl alcohol to prepare tincture of iodine.
However, the solute may be even liquid or gases. All these of miscible liquids (alcohol in water)
are a common example of gas in liquid solutions.


        Sugar       +          water                                  Sugar solution
       (Solute)             (Solvent)




Solution consists of a solvent and one or more solutes.
The substance, which is present in bulk of the solution, is called the solvent and the substance,
which dissolves in the solvent, is called the solute. For example, sugar syrup is sugar dissolved
in water. Here water is solvent and sugar is solute.


Aqueous and Non-aqueous Solutions

The solution formed by dissolving one or more solutes in water is called aqueous solution. For
example, brine (salt solution), vinegar (acetic acid dissolved in water), sugar solution and copper
sulphate solution, ammonium chloride, etc.

And the solution formed by dissolving one or more solute in organic liquids, such as alcohol,
acetone, carbon disulphide are called non-aqueous solutions. For example, sulphur in carbon
disulphide, coal tar in kerosene, nail polish in acetone, kerosene oil in petrol, phosphorus in ethyl
alcohol, iodine in ethyl alcohol, naphthalene in benzene, etc.



Solubility
 The amount of solute required to prepare a saturated solution in a given quantity of solvent at a
given temperature is called the solubility of the solute. It is generally referred to solubility for 100g
of the solvent. Thus, the maximum amount of a solute that can be dissolved in 100 g of a solvent
is called the solubility of that solute in that solvent at a particular temperature and pressure.
Different substances have different solubility in water. For example, the solubility of NaCl in water
               0                                    0
is 46.0g at 25 C and that of KCl is 127.5g at 25 C.


Saturated and unsaturated Solutions

Lets see what are saturated solutions and unsaturated solutions.
To understand what is meant by saturated and unsaturated solutions, perform a simple
experiment.
Take water in a test tube. Add some common salt to it. The salt dissolves. Now, check whether
this solution is saturated or unsaturated. What you have to do is add some more salt into it.
Does the salt dissolve?
If it dissolves, then, it is an unsaturated solution.
If it does not dissolve, then it is a saturated solution.

In your daliy life, you prepare sharbat by dissolving sugar into water. In that if you keep on
dissolving sugar, you might have noticed that, after some time you won’t be able to dissolve more
sugar in it. This indicates solution is saturated completely. The point at which saturation occurs is
known as saturation Point.


An unsaturated solution thus contains less than the maximum possible amount of the solute that
can be dissolved at that particular temperature. A solution is always unsaturated before it
becomes saturated.

Thus,
A saturated solution is one in which, in the presence of excess solute, no more solute can be
dissolved at a particular temperature, i.e. a saturated solution contains the maximum possible
amount of the solute at that temperature

An unsaturated solution is one in which more of the solute can be dissolved at a particular
temperature.




Concentration of Solution


The amount of solute present in the given amount of the solution is called the concentration of a
solution. Concentration of solution can be expressed in many ways, such as mass percent,
volume percent, molarity, etc.

The most common way of expressing concentration of solution is by percentage of solute in the
solution, i.e.,

Mass percent: The concentration of solution as mass percent may be defined as the number of
parts by mass of solute per 100 parts by mass of solution.




                         or
                              Mass of solute
Mass percent (strength) =                       100
                              Mass of solution

And Mass of solution = Mass of solute + Mass of solvent
For example
       Prepare a solution by dissolving 10.0g of sugar in 90.0 of water.

Total mass of solution = mass of solute + mass of solvent
                         = 10.0 g + 90.0 g
                         = 100. 0 g
                          Mass of solute
       Mass percent =                       100
                         Mass of solution
                             10.0
                         =         100 10%
                            100.0
         Thus, mass percent of sugar is 10% in the solution.
Similarly,
A 20 per cent salt solution means that 20 g of salt are present in 100 g of solution.
That is,
Mass of salt (solute) = 20 g
Mass of salt solution = 100 g
Thus, Mass of water (solvent) = Mass of salt solution – Mass of salt
= 100 – 20 = 80 g




Volume percent: The concentration of solution in terms of volume percent may be defined as the
number parts by volume of solute per 100 parts by volume of solution.
                               Volume of solute
Volume per cent (strength) =                      100
                              Volume of solution
For example, prepare a solution by mixing 25 ml of ethyl alcohol in 75 ml of water. Now, total
volume of solution = 25 ml + 75 ml = 100 ml
                    Volume of solute
Volume percent =                       100
                   Volume of solution
            25 ml
                  100  25%
           100 ml
Thus, the volume percent of ethyl alcohol in above solution is 25%




Numerical problem

A sugar solution is prepared by dissolving 20 g of sugar in 250 g of water. What would be the
concentration of this solution?
Solution:
Given,
Mass of sugar (solute) = 20 g
Mass of water (solvent) = 250 g
Mass of sugar solution = Mass of sugar + Mass of water
= 20 + 250 = 270 g




True Solution
In a solution (or true solution), the solute particles cannot be distinguished from the solvent
particles as they are filled into the spaces between the solvent molecules.
The size of the solvent particles is extremely small and cannot be seen even under a microscope.
                                      –8
Their diameter is of the order of 10 cm. True solutions pass through a filter paper without any
residue.

Characteristics of a solution (or true solution)


                                                                 –8
   i.       The size of the solute particles is of the order 10 cm.
  ii.       It is always a homogeneous mixture.
 iii.       The solution gets filtered without any residue.
 iv.        Its particles cannot be seen even with a microscope.
  v.        It does not scatter light.
 vi.        It is quite stable and thus, on standing for a long time, the solute particles do not get
            separated out into a different phase.
 vii.       The properties of solute are retained in the true solution. Thus, a sugar solution is sweet
            in taste and a solution of salt in water is saline in taste.

Suspension


A suspension is a heterogeneous mixture of a solid and a liquid in which the particles of a solid
substance are spread throughout the liquid without dissolving in it. For example, muddy water,
chalk water mixture, paints, etc. In muddy water, sand particles are suspended in water.

Characteristics of a suspension

                                                            –5
   i.       The size of the solute particle is more than 10 cm.
  ii.       It is a heterogeneous mixture.
 iii.       The solute is left as a residue because only the solvent is passed through, on filtration.
 iv.        Its particle can be seen with the naked eye easily.
  v.        It scatters light passing through it.
 vi.        It is unstable and thus, on standing for some time, the particles form a separate phase by
            settling down.

        Colloidal solutions (Colloids)


        In colloidal solution, the size of the particles is intermediate between that of a true solution
        and a suspension. Example: Toothpaste, jelly, starch solution, blood, milk, etc.

        Characteristics of a colloidal solution

                                                                      –7    –5
   i.       The size of the solute particles is of the order of 10 cm to 10 cm.
  ii.       It is a heterogeneous solution, though it may appear to be a homogeneous one.
 iii.       It gets filtered through ordinary filter paper without any residues. They cannot, however,
            pass through parchment membrane.
 iv.        It scatters light passing through it.
  v.        It is quite stable as it does not separate out into different phases even when kept
            undisturbed for a long time.
 vi.        Its particles can be seen through a high-powered microscope.
 vii.       Electrophoresis- The solute particles in a colloidal solution move towards either cathode
            or anode when electric current is passed through it. The movement of colloidal particles
             towards one of the electrodes under the influence of electric field is called
             electrophoresis. Colloidal solutions exhibit electrophoresis due t the presence of charge
             on colloidal particles.
 viii.       Brownian movement of colloids- when colloidal particles are seen under an ultra
             microscope, the particles are found to be in constant motion in zig-zag path in all
             directions. This zig-zag motion of colloidal particles is called Brownian movement. The
             movement of the particles is due to the collisions with the molecules of the dispersion
             medium.




                  Brownian movement of colloidal particles



     Do you know what is delta and how it is formed?
     Delta is a fertile piece of land.
     River water contains charged colloidal particles of clay, sand and other materials. Whereas
     seawater contains salts dissolved in it. When river water comes out into contact with seawater, the
     colloidal particles of river water coagulated. Thus, the level of riverbed rises, due to which
     formation of delta takes place.




              True solutions                 Colloidal solutions                        Suspensions




Distinction between True solutions, Colloidal solutions and Suspensions

Properties         True solutions          Colloidal solutions       Suspensions
                                    –8                       –7                    –5
Size of            Of the order of 10 cm Of the order of 10 cm       More than 10 cm
                                              –5
particle                                 to 10 cm
Nature             Homogeneous             Heterogeneous             Heterogeneous
Visibility         Its particles cannot be Its particles can be seen Its particles can be seen
                seen even under a      through a high-powered      even with the naked eye.
                microscope.            microscope.
Filterability   The solution gets      It gets filtered through    The solute is left as a
                filtered without any   ordinary filter paper       residue because only the
                residue.               without any residues. It    solvent passes through
                                       cannot, however, pass       on filtration.
                                       through a parchment
                                       membrane.
Settling into   Even when kept         It does not separate into   If a suspension is left
a different     standing for a long    its different phases even   standing for some time,
phase           time, the solute       when kept undisturbed       the solid particles form a
                particles do not get   for a long time.            separate phase by
                separated out into a                               settling down.
                different phase.
Scattering of It does not scatter      It scatters light.          It may or may not scatter
light (Tyndall light.                                              light.
effect)
Brownian        It does not show       It shows Brownian           It may show Brownian
movement        Brownian movement      movement                    movement




    Like solution, colloidal solution consists of dispersed phase and dispersed medium.
    Dispersed phase is the substance, which is dispersed, in a large quantity of another
    substance. And dispersion medium is the medium in which these particles are dispersed.

    The colloidal solutions having fluid- like appearance are called sols. The dispersion medium
    in sols is generally liquid. Colloids are sometimes given specific names depending upon the
    nature of the dispersion medium. Some common examples are given below:




    Dispersion Medium                                  Name of colloids


    Water                                              Hydrosols
    Alcohol                                            Alcosols
    Benzene                                            Benzosols
    Gases                                              Aerosols




    Medicines in the colloidal form are easily absorbed by the body tissues and are more effective
    because a colloidal state has large surface area of sol particles and this shows more effective
    absorption.
                         Types of colloidal solutions

Dispersing medium
                          Type                       Examples
+ Dispersed phase
                                                     Mist, Fog, Cloud,
Gas + Liquid              Aerosol
                                                     Insecticide sprays
                                                     Smoke, Dust- storm,
Gas + Solid               Aerosol
                                                     Automobile exhaust
                                                     Froth, Shaving cream,
Liquid + Gas              Foam
                                                     whipped cream
                                                     Milk, Face cream, cod
Liquid + Liquid           Emulsion
                                                     liver oil
                                                     Milk of magnesia,
Liquid + Solid            Sol                        Muddy water, Gold sol,
                                                     Starch sol
                                                     Sponge, Mattress,
Solid + Gas               Solid foam
                                                     Foam, Pumice stone
                                                     Jelly, Cheese, Butter,
Solid + Liquid            Gel
                                                     Shoe polish
                                                     Milk glass, Coloured
Solid + Solid             Solid sol
                                                     gemstone, Some alloys



   Do you know soap solution is a colloidal solution?
   Soap solution is a colloidal solution of sodium stearate in water. It removes dirt by emulsifying the
   greasy matter sticking to cloth.




    Tyndall effect


    When a beam of light passed through a colloidal solution (say a soap solution kept in a glass
    container) in a dark room, the path of the light gets illuminated and becomes visible. This
    happens because the sizes of solute particles are big enough (about 1,000 nm in diameter) to
    scatter the light. This phenomenon of scattering of light by colloidal particles is called Tyndall
    effect.
                                                                     Scattering of light

       Source of light                         True solutions             Colloidal solutions




                                          Tyndall effect



  You must have seen that in forests when light passes through tree, the light becomes visible.
  This is due to scattering of fog (colloidal particles) present in the atmosphere, which is
  nothing but Tyndall effect.




                              Scattering of light through trees




Laws of Chemical Combination
About seven centuries ago, scientists conducted a large number of experiments with a variety of
substances. They carefully observed and documented the results, thus compiling a wealth of
information. Even though their attempts to convert cheaper metals such as copper and lead into
gold failed, they developed different processes such as mixing, separation and distillation of
substances. Antoine L. Lavoisier was the first scientist who by proposing two laws:

Law of conservation of mass that states that mass is neither created nor destroyed in a chemical
reaction This law is also called the law of indestructibility of matter. According to this law
whether a chemical undergoes physical or chemical change the total mass of the products will be
exactly equal to the reactants.

Examples: a) When matter undergoes a physical change - A piece of ice (solid water) is taken in
a small conical flask. It is a well corked and weighed. The flask is now heated gently to melt the
ice (solid) in to water (liquid).

Ice                       Heat                       Water

The flask is again weighed. It is found out that the weight of ice and water is same.

b) When matter undergoes a chemical change- 12g of carbon combines with 32g of oxygen to
form 44g of carbon dioxide.

C          +         O2                          CO 2

12g                  32g                         44g

Solved Example

         1) What weight of silver nitrate react with 1.50g of sodium chloride to produce 2.10g of silver
            chloride and 2.80g of sodium nitrate, if the law of conservation of mass is true?

         Sol. The reaction is

                             AgNO3           +       NaCl          AgCl        +       NaNO 3

                                 X                   1.50g         2.10g               2.80g

         According to law of mass action,

         Mass of reactants = Mass of Products

               X+ 1.50g              =        2.10g+ 2.80g

                      X              =        4.90g – 1.50g

               Mass of NaNO3             =   2.10g

    2.         Study the following reaction
                                           
               NaHCO3  CH3COOH  CH3COONa  H2O  CO2
               If 4.2 of NaHCO3 combines with 10.0 g of CH3COOH to form 12.0 g of CH3COOH & H2).
               Then find out the mass of CO2 formed?
Sol.                       
        NaHCO3  CH2COOH  CH3 COONa  H2O  CO2
           4.2 g        10.0 g                                  x
                                              12.0 g
        According to the law of conservation of mass
        Mass of reactant = Mass of products
        4.2 g + 10.0 g = 12.0 g + x
        x = 4.2 g + 10.0 g – 12.0 g
        x = 14.2 g – 12.0 g = 2.2 g
        Mass of CO2 = 2.2 g




Law of definite proportion which states that in a pure substance the elements are always
present in a definite proportion. Later, quantitative measurements were made. These helped to
develop the Laws of chemical combinations, which led to the concept of atoms being the smallest
independent unit of matter.

The laws of chemical combinations includes:

       Law of Constant Proportions
       Law of Multiple Proportions

Law of Constant Proportions


A pure chemical compound always contains the same elements combined together in a definite
proportion by weight. Whatever may be the source from which a compound is obtained or the
method by which it is prepared, it will always be made up of the same elements in the same fixed
proportion by weight.
For example, water (H2O) obtained from any location or source is always made up of only
hydrogen and oxygen; the proportion is always 1 : 8 (2 : 16) by weight.
Carbon dioxide (CO2) gas obtained by carbon in air or by the action of a dilute acid on a
carbonate is only made up of carbon and oxygen; the proportion of carbon and oxygen is always
3 : 8 (12 : 32). Ammonia, which is a compound of nitrogen and hydrogen, will always combine in
the ratio of 14:3 by mass.




Law of Multiple Proportions


When two elements combine to form two or more compounds, then the weights of one of these
elements, which combine with a fixed weight of the other element, bear a simple ratio to one
another. This law is significant, as in many cases, two elements combine in more than one way,
to form different substances.

For example, in water (H2O), the proportion by weight of hydrogen and oxygen is 2 : 16 or 1 : 8.

And in hydrogen peroxide (H2O2), the proportion by weight of hydrogen and oxygen is 2 : 32 or 1 :
16.
Hence, the ratio of oxygen combining with a fixed amount of hydrogen (say 1 g) is 8 g : 16 g or
1 : 2, between water and hydrogen peroxide.
Similarly, in carbon monoxide (CO), the proportion by weight of carbon and oxygen combining is
12 : 16 or 3 : 4.

And in carbon dioxide (CO2), the proportion by weight of carbon and oxygen combining is 12 : 32
or 3 : 8.

Hence, the ratio of oxygen combining with a fixed amount of carbon (say 3 g) is 4 g : 8 g or 1 : 2,
between carbon monoxide and carbon dioxide.

One more interesting example is of Sulphur dioxide and sulphur trioxide.

In sulphur dioxide (SO2), the proportion by weight of sulphur and oxygen combining is 32 : 32 or 1
: 1.

And in sulphur trioxide (SO3), the proportion by weight of sulphur and oxygen combining is 32: 48
or 2 : 3.

Hence, the ratio of oxygen combining with a fixed amount of sulphur is 1 : 3, between sulphur
dioxide and sulphur trioxide.




Atomic Theory

About 400-500 BC, Greek philosopher Leuappus and Democritus suggested that if we go on
dividing the matter, a stage will come when particles obtained cannot be divided further similar to
the Kanada’s theory of parmanus.

However, not much experimental work could be done until Lavoisier gave the law of conservation
of mass and law of constant proportions. Guided by the laws of chemical combination and other
facts, in early 18th century, John Dalton (1766-1844) proposed the basic theory about the nature
of the matter that all the matters (element, compound or mixture) are made of tiny particles
invisible to naked eyes called atoms.

Dalton’s Atomic Theory of Matter

Dalton proposed his atomic theory as:

   i.   All matters are made up of very tiny particles, called atoms.
  ii.   Atoms are indivisible. (Tomio in Greek means to divide or break; so, ‘atomio’ means
        indivisible)
 iii.   Atoms can neither be created nor be destroyed.
 iv.    Atoms of the same element are identical in all respects (i.e. they have the same physical
        and chemical properties).
  v.    Atoms of different elements are different in all respects (i.e. they have different physical
        and chemical properties).
 vi.    Atoms of different elements combine in a simple whole number ratio to form compounds.
 vii.   Atoms of same element can combine in more than one ratio with another element to form
        two or more compounds. For example, carbon and oxygen combine to give carbon
        monoxide and carbon dioxide. In carbon monoxide, one atom of carbon combine with
        one atom of oxygen while in carbon dioxide, two atoms of oxygen combines with one
        atom of carbon.
We can summarize from the above theory that atoms of different particles are different. An
atom is the smallest particle that maintains its chemical properties in all the chemical and
physical changes.




Merits of Dalton’s Atomic Theory


The Dalton’s atomic theory helped the scientist to have a clear picture of matter. According to
the atomic theory, an atom is the smallest particle of an element, which may or may not have
independent existence

The Dalton’s atomic theory gives a simple explanation for:

a. Formation of molecules
b. The laws of chemical combinations

Demerits of Dalton’s Atomic Theory


The demerits of Dalton’s theory are:

a. It states that atoms are indivisible; but it was later discovered that atoms have sub-atomic
   particles — protons, neutrons and electrons.
b. It could not explain why atoms of the same element always do not have the same mass.
   These atoms are called isotopes.
c. It states that atoms of different elements have different masses. But, now it is known that
   even atoms of different elements can have the same mass.
d. It could not explain the difference in properties of various materials made up of the same
   type of atoms. For example, charcoal and diamond have different properties though they
   are made up of only carbon.




Today modern technology enables to take photographs of atoms. With the help of Scanning
Tunneling Microscope (STM), image of the surface of atom can be produced.




                       A detail photographic image taken by STM

An atom is the smallest particle of an element that can take part in a chemical reaction.
Atoms and Molecules


Only the atoms of noble gases like helium, neon, argon, krypton, etc., are unreactive and
exist in the free state.
Atoms combine to form a molecule.

A molecule is the smallest particle of a substance (element or compound) that can exist in
the free state under normal conditions and shows all the properties of that substance.

       Molecules of elements — hydrogen (H2), nitrogen (N2), oxygen (O2).
       Molecules of compounds — water (H2O), sodium chloride (NaCl), methane (CH4).

Thus, a molecule contains two or more atoms. For example, a molecule of water (H2O)
contains 2 atoms of hydrogen and 1 atom of oxygen. Similarly, a molecule of carbon dioxide
(CO2) contains one atom of carbon and two atoms of oxygen.

                    Atomic Masses of some Common elements


                 Element          Symbol        Atomic mass or Atomic weight
             Hydrogen               H                        1
             Carbon                 C                       12
             Nitrogen               N                       14
             Oxygen                 O                       16
             Sodium                 Na                      23
             Magnesium              Mg                      24
             Aluminium              Al                      27
             Phosphorus             P                       31
             Sulphur                S                       32
             Chlorine               Cl                     35.5
             Potassium              K                       39.
             Calcium                Ca                      40.
             Iron                   Fe                      56
             Copper                 Cu                     63.5




Molecules can be classified into two types:

(1) Homoatomic molecules – These molecules are made up of two or more atoms of the
    same element. Theses are called elementary molecules. For example, hydrogen gas
    consists of two atoms of hydrogen (H2), Simlarly, oxygen (O2) gas is consist of two atoms
    of oxygen.Such gases are called diatomic gases because these consist of two atoms of
    the same element. The number of atoms in amolecule of an element is called atomicity.
    Thus the molecules may be monoatomic, diatomic and triatomic etc. containing one ,two,
    three, four atoms respectively. For example:

Monoatomic- a molecule containing one atom like Helium(He), Neon(Ne), argon (ar).

Diatomic- a molecule containing two atoms like oxygen (O2), hydrogen (H2), nitrogen (N2).

Triatomic- a molecule containing three atoms like ozone(O3).
(2) Heteroatomic molecules- These molecules are made up of two or more atoms of the
    different elements. Theses are called compound molecules. For example, a molecule of
    water (H2O) contains 2 atoms of hydrogen and 1 atom of oxygen. Similarly, a molecule of
    carbon dioxide (CO2) contains one atom of carbon and two atoms of oxygen.




Monoatomic         Diatomic            Triatomic
molecules of He    molecules of Cl2    molecules of O3




Molecules of HCl Molecules of H2O      Molecules of H2O2

Solved Example

Classify the following as homoatomic and heteroatomic molecules:

CO2, O3, CO, N2, PCl5

(a) CO2 = It consists of C and O atoms. Therefore, it is heteroatomic molecule.

(b) O3 = It contains only O atoms. Therefore, it is homoatomic molecule.

(c) CO = It is made ur of C and O atoms. Therefore, it is heteroatomic molecule.

(d) N2 = It contains only N atoms. Therefore, it is homoatomic molecule.

(e) PCl5 = It consists of P and Cl atoms. Therefore, it is heteroatomic molecule.




Chemical formulae


Any molecule can be represented by the following chemical formulae:
i.        Molecular formula: A molecular formula of a substance makes it clear that how many
          atoms of each kind of element are present in one molecule. For example, H 2O, CO2, CH4,
          etc.
ii.       Structural formula: A structural formula is a chemical formula represents how the atoms
          are bonded with each other in a molecule. For example, water can be represented as H –
          O – H.

      Some properties of elements and compounds depend upon the arrangement of the atoms
      and molecules.

      Atomic Mass


      The main feature of Dalton’s theory was that he proposed that each atom of the element has
      certain mass known as atomic mass and atoms of different elements combine in a simple
      whole number ratio to form compounds. Atoms of same element can combine in more than
      one ratio with another to form two or more compounds. Scientist considered various atomic

      mass units (amu) but at last they took     of the mass of an atom of naturally occurring
      oxygen as oxygen reacts with most of the elements to form various compounds and also, this
      way of calculating mass of different atoms usually gave mass in whole number.
      But, all the atoms of oxygen do not have same atomic mass. Oxygen found in nature is a
      composition of a mixture of atoms of different masses. So, the relative proportion of these
      different atoms (isotopes) could change with time and place. To avoid this problem,

      nowadays the atomic mass unit is defined as     of the mass of one atom of a particular
      isotope of the carbon (C–12).
      The actual mass of one atom of this isotope of the carbon (C–12) has a mass 1.992         g.

      Gram atomic mass is atomic mass of an element expressed in grams.




      Molecular Mass


      The molecular mass of a substance is the sum of the atomic masses of all the constituent
      atoms of the molecule. So, the average mass of the molecule of a substance is expressed in
      amu.
      For example, molecular mass of carbon dioxide         can be calculated as:
      Molecular mass of       = Atomic mass of C + 2 (Atomic mass of O)
      = 12 +2 (16)
      = 12 + 32
      = 44 amu
      Similarly, molecular mass of NaCl is = Atomic mass of Na + Atomic mass of Cl
      = 23 + 35.5
      = 58.5 amu

      Numerical examples

      Calculate the molecular mass of the following
(a)       PCl5 (Phosphorus Penta chloride)
          (Atomic masses are: P = 31, Cl = 35.5)
          Solution:
          1 × At. Mass of P + 5 × At. Mass of Cl
          1 × 31 + 5 × 35.5 = 31 + 177.5
          = 208.5 a.m.u

(b)       NH3 (Ammonia)
          (Atomic masses are: N = 14, H = 1)
          Solution:
          1 × At. Mass of N + 3 × At. Mass of H
          = 1 × 14 + 3 × 1
          = 14 + 3
          = 17 a.m.u


(c)      CaCO3 (Calcium Carbonate)
         (At. Masses are: Ca = 40, C = 12, O = 16)
         Solution:
         = 1 × At. Mass of Ca + 1 × At. Mass of C + 3 × At. Mass of O
         = 1 × 40 + 1 × 12 + 3 × 16
         = 40 + 12 + 48
         = 100 a.m.u

(d)      HNO3 (Nitric acid)
         (Atomic masses are: H = 1, N = 14, O = 16)
         Solution:
         1 × At. Mass of H + 1× At. Mass of N+ 3× At. Mass of O
         1 × 1 + 1 × 14 + 3 x 16 = 1 + 14 + 48
         = 63 a.m.u




Gram molecular mass is molecular mass of a substance expressed in grams.




      Mole Concept

      Whenever you have to buy egg or oranges, you always do it in the form of dozen. Dozen is a
      counting unit, which depicts the quantity of 12.

      Thus, Mole is the basic counting unit (like dozen (12), ream (20) or gross (144)) for atoms,
      molecules, ions, electrons, etc. A mole represents              particles irrespective of their
      nature.


                                             1 mole



                    6.023x1023             6.023x1023               6.023x1023
                    Particles              Atoms                    Molecules
This number                 is called Avogadro’s number and is represented by N o or   or L.
A mole is also related to the mass of the substance. Thus,
1 mole (              ) of atoms of an element = Gram atomic mass of the element
1 mole (              ) of molecules of a substance = Gram molecular mass of the
substance

For example:

1 mole of oxygen atoms = 16 g
or 2 moles of oxygen atoms = 2 × 16
= 32 g
1 mole of CO2 molecule = 44 g
or 2 moles of CO2 molecules = 2 × 44
= 88 g
In terms of the unit for the amount of substance, mole is defined as follows,

A mole is that amount of the substance which contains as many particles (atoms, molecules,
ions, etc.) as there are atoms in 12 g of C-12.


The number of atoms present in 12 g of C–12 =




Numerical examples

        1. What is the mass in grams of 5 moles of nitrogen atoms?
           Solution:
           1 mole of nitrogen atoms = Gram atomic mass of nitrogen atom
           i.e. 1 mole of nitrogen atoms = 14 g
           Thus, 5 moles of nitrogen atoms = 14 × 5
           = 70 g




        2. Calculate the number of moles in 34 g of NH3
           Solution:
           1 mole of NH3 = Gram molecular mass of NH3
           Gram molecular mass of NH3 = Gram atomic mass of N + 3 (Gram atomic mass
           of H)
           = 14 + 3 (1)
           = 14 + 3 = 17 g
           Thus, 1 mole of NH3 = 17 g
        or 17 g of NH3 = 1 mole



        Thus, 34 g of NH3 =




    3. If one mole of water molecules weighs 18 g, calculate the mass of one molecule
       of water in grams.
       Solution:
       Mass of one mole of water = 18 g
        i.e. Mass of                     water molecules = 18 g



        Thus, Mass of 1 water molecule




    4. What is the mass of 1 atom of carbon?
       Solution:




1 mole of carbon                = 12 grams
                                                  23
                                = 6. 023     10 atoms of it
                        23
i.e. mass of 6.023     10
atoms of C                      = 12 g


  mass of 1 atom of C

                                             -23
                                = 1.99     10      g

    5. What is the weight in grams of 0.5 mole of CO2?
       Solution

Molecular mass
                     = (1    12) + (2       16)
of CO2
                            C                          O
                     = 12 + 32 = 44
  1 Mole of CO2      = 44 gms
 0.5 mole of
                     = 0.5      44 = 22 gms
CO2

    6. How many molecules of water are present in 10 g of it?
       Solution
        Molecular weight of H2O           = (2    1) + (1       16)
             1 mole of H2O                = 18 g.
                                                           23
                                          = 6.023        10 molecules
                                                           23
        No. of molecules present in = 6.023              10
        18 g of water
        No. of molecules in 10 g of
                                          = 0.5     44 = 22 gms
        water




                                                           22
                                          = 33.46        10 molecules

              7. Calculate the number of moles in 34 g of NH3.
                 Solution:

        Molecular mass
                             = (1   14) + (3        1)
        of NH3
                                      N                           H
                             = 14 + 3 = 17
             17 g of NH3     = 1 mole

             34 g of NH3     = 2 mole




Percentage composition


If we know the formula of a compound, then we can calculate the percent of any constituent
element.
Thus, percentage composition of a compound corresponds to the mass of each element in
100 g of the compound.

Let us take a compound XY made of elements X and Y.




Similarly,




Numerical example:
1. Calculate the percentage composition of hydrogen and oxygen in water, H2O?
Solution:
Molecular mass of H2O = 2 (Atomic mass of H) + Atomic mass of O
= 2 (1) + 16 = 18 g
Mass of H in H2O = 2 (Atomic mass of H)
= 2 (1) = 2 g




Mass of O in H2O = Atomic mass of O
= 16 g




Thus, percentage composition of hydrogen and oxygen in water is 11.11 % and 88.89 %.


2.Calculate the % of nitrogen in urea (H2N)2CO
Solution:
Molecular mass of urea = 2 (2 + 14) + 12 + 16 = 60
Mass of 1 mole of urea = 60 g
1 mole of urea contains 2 moles of N(28 g)
                          28
       % of N in urea =      100  46.7%
                          60



Empirical and Molecular Formulae



Empirical formula of a compound is the simplest formula, which gives the simplest ratio in
whole number of atoms of different elements present in one molecule of the compound.

Molecular formula of a compound is the formula, which gives the actual number of atoms of
different elements present in one molecule of the compound.



For example,


                                Molecular formula                Empirical formula


Hydrogen Peroxide               H2O2                             HO
Glucose                          C6H12O6                          CH2O

Acetic acid                      CH3COOH                          CH2O

Phosphorus pentoxide             P2O5                             P2O5




Empirical Formula Mass

E.F.M. of a substance is equal to the sum of atomic masses of all atoms in the empirical formula
of the substance. Molecular formula is whole number multiple of empirical formula.
Thus,
         Molecular Formula = empirical formula × n
               Molecular formula         Molecular Mass
          n                       
               Empirical formula     Empirical formula Mass
For example,
         The empirical formula of the compound Hydrogen peroxide (H2O2) is HO which shows
         that H and O are present in the simplest ratio of 1 : 1



Determination of the empirical formula of a compound

To determine the empirical formula of a compound some steps are given below:

1. Divide the percentage of each element in the compound by its atomic mass. This gives the
atomic ratio of the various elements in the molecule of the compound.




If the total percentage of all the elements in the compound detected by chemical analysis does
not come out to be 100%, then oxygen is assumed to be present in a quantity that brings the total
percentage to 100%.

For example, a compound on chemical analysis showed a percentage composition of 92.4%
carbon and 7.6% hydrogen.

Total percentage of carbon (C) and hydrogen (H) = 92.4 + 7.6 = 100%
Therefore, Oxygen is absent




    2. Divide the atomic ratio by the smallest value among them to get the simplest ratio of the
       elements.
       For example, the smallest value among the atomic ratio for carbon (C) and hydrogen (H)
       in the above example is 7.6.




   3. Make the simplest ratio of the elements to their nearest whole number. (If necessary, this
      can be done by multiplying with a suitable integer.) This is known as the simplest whole
      number ratio,
      For example, in the above example,
      Simplest whole number ratio for C = 1
      Simplest whole number ratio of H = 1
   4. The ratio of the simplest whole number ratio for the various elements in the compound
      gives the ratio of the number of atoms of the elements present in the compound.
      For example, in the above example
      Simplest whole number ratio for C = 1
      Simplest whole number ratio of H = 1
      Therefore,
      Simplest whole number ratio of C : Simplest whole number ratio of H :: 1 : 1
      Thus, the number of atoms of carbon (C) and hydrogen (H) present in one molecule of
      compounds is given by the ratio 1 : 1
   5. Write the symbols of the various elements in the molecules side by side followed by the
      numerical value (ratio of the number of atoms of that element present) at the right lower
      corner of each symbol. This gives the empirical formula of the compound.
      For example, in the above example, the empirical formula is given as C1H1, i.e. CH
      Lets see some more examples



Numerical examples:


S. Percentage Percentage Atomic Atomic             Simplest Simplest Ratio of Empirical
No. composition of oxygen, mass of ratio           ratio    whole the         formula
    by chemical if present the                              number number
    analysis               element                          ratio    of
                                                                     atoms
                                                                     present
                                                                     in the
                                                                     element
1. Copper       Oxygen       63.5                           1
   (47.3%)      (100 –
                {47.3 +                                                 Cu:Cl ::
                                                                                   CuCl2
                52.7})                                       2           1:2
   Chlorine     = 0%
   (52.7%)                   35.5
2. Carbon       Oxygen       12                              2×2=4
   (57.8%)      (100 –
                {57.8 +
   Hydrogen     3.6}         1                               1.5 × 2   C:H:O ::
                                                                                C4H3O2
   (3.6%)       = 38.6% )                                    =3        4:3:2

                             16
                                                             1×2=2
3. Sodium           Oxygen         23                           1×2=2
   (29.11%)         (100 –
                    {29.11 +
                                                                          Na:S:O
   Sulphur          40.51}         32                           1×2=2
                                                                          :: 2 : 2 : Na2S2O3
   (40.51%)         = 30.38%
                                                                              3
                                   16                           1.5 × 2
                                                                =3



Q. A compound is found by analysis to consist of 54.54% of carbon, 9.009% of hydrogen and
        36.36% of oxygen. The vapour density of the compound was found to be 44. Find the
        molecular formula of the compound
       Solution:
        Calculate of empirical formula

          Element    %         Atomic mass    Relative number    Simplest ratio   Simples whole
                                              of moles           of moles         no. ratio
                                              54.54              4.53
          C          54.54     12                     4.53           2          2
                                               9.09              2.27
                                              9.09               9.09
          H          9.09      1                     9.09            4          4
                                                 1               2.27
          O          36.36     16             36.36              2.27
                                                      2.27           1          1
                                                16               2.27


             Empirical formula is C2H4O
              Calculated of molecular formula:
       Empirical formula mass = 12 × 2 + 1 × 4 + 16 × 1 = 44
       Molecular mass = 2 × vapour density = 2 × 44 = 88
                 Molecular Mass         88
       n                                 2
            Empirical Formula Mass      44
       Molecular formula = Empirical Formula × n
                         = C2H4O × 2
                         = C4H8O2

Formula
                               Mass of solute
                % solute =                      100
                               Mass of solution
                                               Volume of solute
                Volume per cent (strength) =                     100
                                              Volume of solution
                Mass of reactants = Mass of Products
                1 mole (               ) of atoms of an element = Gram atomic mass of the
                 element


                 1 mole (              ) of molecules of a substance = Gram molecular mass of
              the substance
                                               Mass of an atom
              Atomic mass =
                                                                  12 C
                                    th
                                1
                                         mass of a carbon atom
                               12
                                                Mass of a molecular
              Molecular mass =
                                                                       12 C
                                         th
                                     1
                                              mass of a carbon atom
                                    12


       

                                  1 gram abonic mass
              Mass of 1 atom =
                                          6.022  1023
                                              Mass of the atoms
      Mass percentage of element =                              100
                                               Molecular Mass

              Molecular formula = (empirical Formula)n
                              Molecular mass
                       n
                           Empirical Formula Mass




Points to remember


      Matter is a physical object or material that has mass and occupies space.
      Matter can be classified as solid, liquid and gas.
      The bonding forces between the basic particles (atoms/ molecules) of the matter of which
       it is made of, is known as intermolecular force.
      Evaporation is a process in which the liquid disappears in the gaseous state into the
       environment.
      On heating, the solids change into the liquid and liquid into the gaseous state.
      On applying pressure, solids cannot be compressed, liquids can be slightly compressed
       and gases can be easily compressed.
      Any homogeneous mixture of two or more substances (solid, gas or liquid) is a solution.
      The substance that gets dissolved to form a solution is known as solute.
      The liquid in which the solute gets dissolved is known as solvent.
      Water is called the universal solvent.
      The solution formed by dissolving different substances in water is called aqueous solution
       and if the solvent is any organic liquid, then it is non-aqueous solution.
      In a true solution, the solute particles are present as molecules or ions giving a
       homogeneous mixture of solute and solvent forming a single phase. If the size of solute
       particles is bigger than true solution and smaller than those of suspension, then it is
       colloidal solution and in case when the small solid particles are spread throughout a liquid
       without dissolving in it, then it is known as suspension.
      The phenomenon of scattering of light by colloidal particles is called Tyndall effect.
      Unsaturated solution is a solution in which more quantity of solid can be dissolved without
       raising the temperature while saturated solution is a solution in which no more solute can
       be dissolved at the same temperature.
      The amount of solute present in the given mass or volume of the solution is called the
       concentration of a solution.
         The maximum amount of a solute that can be dissolved in 100 mg of a solvent is
          solubility of that solute in that solvent at a particular temperature and pressure
         Law of conservation of mass that states that mass is neither created nor destroyed in a
          chemical reaction.
         Law of definite proportion which states that in a pure substance the elements are always
          present in a definite proportion.
         All matters are made of atoms.
         Atoms of different elements combine in a simple whole number ratio to form compounds.
         A molecule is the smallest particle of a substance (element or compound) that can exist
          in the free state under normal conditions and shows all the properties of that substance.
         A molecular formula of a substance makes it clear that how many atoms of each kind of
          element are present in one molecule.
         A structural formula is a chemical formula that represents how the atoms are bonded with
          each other in a molecule.

         Atomic mass unit is defined as    of the mass of one atom of a particular isotope of the
          carbon (C–12).
         The actual mass of one atom of this isotope of the carbon (C–12) has a mass 1.992
             –23
          10 g.
         The molecular mass of a substance is the sum of the atomic masses of all the constituent
          atoms of the substance.
         The formula mass of any compound whether molecule or not is the sum of atomic
          masses of all constituent atoms.
                                                                      23
         The number of atoms present in 12 g of C-12 = 6.0223 10 .
                                    23
          This number 6.0223 10 .is called Avogadro's constant and is represented by NA.
         If we know the formula of a compound, then we can calculate the mass percentage of
          any constituent element.




i.        The empirical formula of a compound is a chemical formula that shows the relative
          number of atoms of each element.
ii.       The molecular formula of a compound is a chemical formula that shows the actual
          number of different types of atoms present in one molecule of the compound.
                                                                                           Dalton’s Atomic
                                            Matter                                         Theory



                     Classification of Matter
                                     Chemical classification         Laws of chemical combination
               Physical

                                                                                Law of Constant
                                                                                Proportions
  Solids    Liquid        Gases         Pure substances             Mixtures
                                                                                       Law of Multiple
                                                                                       Proportions

               element                     compounds          Homogeneous      Heterogeneous
                                                              mixture


metal      non-metal       metalloids




                                            1 Mole

                                         6.023 × 1023
                                           particles




                            1 Gram                           1 Gram
                          atomic mass                     molecular mass


                     H             O                    H2O        CO2

                     1g            16g                  18g         44g

                                  Mole Concept
Practice Exercise

Very Short Answer Type Questions

         1.    Define matter.
         2.    Name the three states of matter.
         3.    Name the Indian philosopher who proposed the theory of matter.
                                                        0
         4.    What is the physical state of water at 0 C?
         5.    What is the general name of the substances which contain atleast two pure
               substances and show the properties of their constituents?
         6.    Define intermolecular force.
         7.    In which of the following substances, the intermolecular force is the strongest and
               in which the intermolecular force is the weakest?
               Water, sodium chloride, carbon dioxide, alcohol
         8.    Among the three states of matter which shows the maximum movement of
               molecules?
         9.    Name two factors which decide the state of a substance?
         10.   A matter when applied pressure neither changes the shape nor distorts. What is
               the state of matter?
         11.   Why solids have high density in comparison to liquids and gases?
         12.   What is lattice?
         13.   Classify the following into elements and compounds:
               Tungsten, Graphite, Iron sulphide, Copper, Washing soda
         14.   Which type of mixture is brass?
         15.   What is solution?
         16.   What is an unsaturated solution?
         17.   A substance is dissolved in the alcohol. What can you say about this solution?
         18.   What should be the size of particles for solution to be a suspension?
         19.    How is the dispersed phase different from the dispersed medium?
         20.   What is the particle size range of colloids?
         21.   Give two examples of gels.
         22.   Give an example of sol and aerosol.
         23.   Which of the two will scatter light- muddy water or starch solution? Why?
         24.   What is the importance of mass percentage?
         25.   What happens to the solubility of gases when temperature increases?
         26.   Who proposed the modern atomic theory?
         27.   Define formula mass.
         28.   How many atoms are present in one mole of atoms?
         29.   What is the mass in grams of 2.5 moles of Cl? (atomic mass of Cl=35.5)
         30.   A compound on chemical analysis showed a percentage composition of 80.4%
               carbon and 8.5% hydrogen. Then the percentage composition of oxygen in this
               compound is?
         31.   What is Avogadro’s number?
         32.   25 g of sugar is present in 100 g of sugar solution. What is the amount of water
               present in this solution?
         33.   A compound on chemical analysis showed a percentage composition of 80.4%
               carbon and 8.5% hydrogen. Then the percentage composition of oxygen in this
               compound is?
         34.   Calculate the molecular mass of Na2CO3. The relative atomic masses are Na =
               23, C = 12, O = 16
         35.   Which element is used as reference for atomic mass scale?
         36.   If C6H12O6 is the molecular formula of glucose, what could be its empirical
               formula?
         37.   What is the percentage of oxygen in HgO? (Atomic mass of Hg is 200 and that of
               O is 16).
Short Answer Type-I Questions

  1. Give reason why solid and liquid can’t be compressed, whereas a gas can be easily
        compressed.
  2.    What is the significance of symbol?
  3.    Why air is considered a mixture and not a compound?
  4.    Differentiate between an element and a compound.
  5.    Give two examples of each of the following.
        Solid dissolved in the liquid
        Liquid dissolved in the liquid
        Gases dissolved in the liquid
        Solid dissolved in the solid
  6.    Discuss effect of temperature on solids.
  7.    Discuss the effect of pressure on gases.
  8.    Differentiate between effusion and diffusion.
  9.    Define the terms solute and solvent. Give an example for each.
  10.   Define concentration of the solution.
  11.    What is the effect of temperature on solubility?
  12.   Give two examples of each of true solution and colloidal solution.
  13.   What is Tyndall effect?
  14.   What does law of conservation of mass states? Give an example also.
  15.   State law of the multiple proportions.
  16.   Define atomic mass of an element. Give the atomic mass for nitrogen and oxygen.


  17. Why scientists preferred to calculate the atomic mass as     of the mass of an atom of
        naturally occurring oxygen?

  18. Define molecular mass. What is the unit of molecular mass?
  19. What is molecular mass of
            (a)C6H12O6
             (b)C12H22O11
  20.   Define mole. How many atoms are there in 1 mole of atoms?
  21.   Which contains more atoms, 50g of Na or 50g of K.?
  22.   {Atomic mass of C = 12, H = 1, O = 16 a.m.u.}Calculate the mass of 5.2 moles of iodic
        acid (HIO3).
        [Atomic mass of H = 1, I = 127, O = 16]




Short Answer Type-II Questions

  1.    Determine the proportion of elements in the following compounds:
  2.    Write down the properties of liquid state.
  3.    Write down the properties of gaseous state.
  4.    Write down the properties of solid state.
  5.    By giving suitable example reasons explain why water is a compound not mixture.
  6.    Differentiate between true solution, colloidal solution and suspension?
  7. Write down the characteristics of the true solution.
  8. What are three major properties shown by colloidal solution?
  9. What happens when an electric current is passed through a colloidal solution? Explain
       the phenomena associated with this.
  10. What is the mass in grams of each of the following?
           (a) 0.2 mole of Na
           (b) 1.5 mole of Cu
           (c) 0.5 mole of K
  11.  Calculate the number of molecules in 93 g of salt (NaCl).
  12. If in a certain amount of sulphur dioxide (    ) the number of molecules is 2, what is the
       weight of sulphur dioxide (    )?
  13. Calculate the percentage composition of CaCO3
  14. A compound ‘X’ after analysis gave 75% of C, 25% of H. What is its empirical formula?
  15. Calculate the number of molecules in 85 g of ammonia (     ).
  16. In carbon monoxide, 12 g of carbon combine with 16 g of oxygen. And in carbon dioxide,
       12 g of carbon combine with 32 g of oxygen. Is this information is in accordance with the
       law of multiple proportion?



Long Answer Type Questions. (5 marks each)


  1. Explain the states of matter according to the intermolecular force between constituent
       particles.
  2.   Make a comparison between true solution colloidal solution and suspension.
  3.   State Dalton’s atomic theory. Why this theory failed to explain atoms completely?
  4.   A compound of boron and hydrogen was found to have 78.2% boron and 21.8%
       hydrogen. Its molecular mass has been found to be 27.6. What is the molecular formula
       of the compound?(atomic masses: B=10.8, H=1 )
  5.   A compound contains 85.7% C and 14.3% H by mass. The molecular mass of the
       compound is 28. Calculate the molecular formula of the compound. (atomic masses:
       C=12, H=1 )
                                  Chapter Test- Nature of Matter



Time: 1 hour                                    Marks: 20

Instructions:
    1. All questions are compulsory.
    2. Marks for each question are given against it.
    3. Answers of 1 mark questions have to be answered in one word or maximum in one
        sentence.
    4. Answers of 2 marks questions may not normally exceed 40 words each.
    5. Answers of 3 marks questions may not normally exceed 60 words each.
    6. Answers of 5 marks questions should not be more than 100 words each.


   1. Classify the following into elements and compounds:
   Iron, water, copper, iron sulphide, graphite, baking soda (1 mark)

   Solution:
   Element = Iron, copper, graphite (1/2 mark)
   Compound= water, iron sulphide, baking soda(1/2 mark)


   2. Calculate the molecular mass of CHCl3 (Chloroform). The relative atomic masses are: C =
   12, H = 1, Cl = 35.5 (1 mark)

   Solution:
       = 1 × At. Mass of C + 2 × At. Mass of H + 2 × atomic mass of Cl
       = 1 × 12 + 2 × 1 + 2 × 35.5
       = 12 + 2 + 71
       = 85 a.m.u

   3.What is the mass of 2 moles of O? (1 mark)

   Solution:
     Mass of 2 moles atoms of oxygen
     = 0.2 moles x At. Mass of oxygen
     = 0.2 moles x 16 g/mol.
     = 3.2 g

   4.State the law of multiple proportions. (2 marks)

   Solution:
     When two elements combine to form two or more compounds, then the weights of one of
   these elements, which combine with a fixed weight of the other element, bear a simple ratio
   to one another. (1 mark)

   For example, in water (H2O), the proportion by weight of hydrogen and oxygen is 2 : 16 or 1 :
   8.

   And in hydrogen peroxide (H2O2), the proportion by weight of hydrogen and oxygen is 2 : 32
   or 1 : 16.
   Hence, the ratio of oxygen combining with a fixed amount of hydrogen (say 1 g) is 8 g : 16 g
   or
   1 : 2, between water and hydrogen peroxide. (1 mark)
5. Calculate the number of molecules of sulphur (S8) in 16 g of solid sulphur. (2 marks)
Sol.    Molar mass of S8 = 8 × 32 g/mol
                          = 256 g/mol
        Mass of solid sulphur = 16 g
                                      16 g
        Number of moles of S8 =               = 0.0625 mol(1 mark)
                                  256 g / mol
                                                        23
        Number of molecules = 0.0625 ml × 6.022 × 10 molecular/mole
                                         22
                             = 3.76 × 10 molecular(1 mark)

     6. What is dispersed phase and dispersion medium? (2 marks)

     Solution:
     Dispersed phase is the substance, which is dispersed, in a large quantity of another
     substance. And dispersion medium is the medium in which these particles are dispersed.


     7. 6. Define:
     (a) A mole
     (b) Aqueous solution
     (c) Unsaturated solution
     (3 marks)

     Solution:
     (a) A mole = A mole is that amount of the substance which contains as many particles
         (atoms, molecules, ions, etc.) as there are atoms in 12 g of C-12. (1 mark)

     (b) Aqueous solution = The solution formed by dissolving one or more solutes in water is
         called aqueous solution. For example, brine (salt solution), vinegar (acetic acid dissolved
         in water), sugar solution, etc. (1 mark)


     (c) Unsaturated solution= An unsaturated solution is one in which more of the solute can be
         dissolved at a particular temperature. (1 mark)




     8.Give differences between true solution, suspension and colloidal solution in terms of:
         i. Particle size
         ii. Visibility of particles
         iii. Filterability
     (3 marks)

     Solution:
Properties True solutions           Colloidal solutions             Suspensions
                                                      –7                         –5
Size of         Of the order of     Of the order of 10 cm to        More than 10 cm
                  –8                  –5
particle        10 cm               10 cm
Visibility      Its particles       Its particles can be seen       Its particles can be seen
                cannot be seen      through a high-powered          even with the naked eye.
                even under a        microscope.
                microscope.
Filterability   The solution gets   It gets filtered through        The solute is left as a
                filtered without    ordinary filter paper without   residue because only the
         any residue.        any residues. It cannot,      solvent passes through on
                             however, pass through a       filtration.
                             parchment membrane.



9. A compound is found by analysis to consist of 54.54 % of carbon, 9.09 % of hydrogen and
36.36 % of oxygen. The molecular mass of the compound was found to be 88. Find the
molecular formula of the compound. (5 marks)

  Solution:
   Calculate of empirical formula

     Element   %        Atomic mass    Relative number      Simplest ratio   Simples whole
                                       of moles             of moles         no. ratio
                                       54.54                4.53
     C         54.54    12                     4.53             2          2
                                        9.09                2.27
                                       9.09                 9.09
     H         9.09     1                     9.09              4          4
                                          1                 2.27
     O         36.36    16             36.36                2.27
                                               2.27             1          1
                                         16                 2.27


         Empirical formula is C2H4O
          Calculated of molecular formula:
   Empirical formula mass = 12 × 2 + 1 × 4 + 16 × 1 = 44
   Molecular mass = 88
             Molecular Mass         88
   n                                 2
        Empirical Formula Mass      44
   Molecular formula = Empirical Formula × n
                     = C2H4O × 2
                     = C4H8O2

				
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