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Science can be viewed as a continuing human effort to systematise knowledge for describing and understanding nature. For the sake of convenience, science is sub-divided into various disciplines: chemistry, physics, biology, etc. The branch of science that studies the preparation, properties, structure and reactions of material substances is called chemistry.

Chemistry deals with the composition, structure, properties and interactions of matter and is of much use to human beings in daily life. These aspects can be best described and understood in terms of basic constituents of matter are atoms and molecules. That is why chemistry is also called the science of atoms and molecules. 


Anything which has mass and occupies space is called matter. Everything around us, for example, book, pen, pencil, water, air, all living beings, etc., are composed of matter. You know that they have mass and they occupy space. 

 States of Matter


Particles are held very close to each other in solids in an orderly fashion and there is not much freedom of movement. In liquids, the particles are close to each other but they can move around. However, in gases, the particles are far apart as compared to those present in solid or liquid states and their movement is easy and fast. 

Arrangement of particles in solid, liquid and gaseous state
(Image credit: NCERT)

Because of such arrangement of particles, different states of matter exhibit the following characteristics 
(i) Solids have definite volume and definite shape. 
(ii) Liquids have definite volume but do not have a definite shape. They take the shape of the container in which they are placed.
(iii) Gases have neither definite volume nor definite shape. They completely occupy the space in the container in which they are placed.

These three states of matter are interconvertible by changing the conditions of temperature and pressure

Classification of Matter

Classification of matter (Image credit: NCERT)

At the macroscopic or bulk level, matter can be classified as mixture or pure substance. These can be further sub-divided 


A mixture contains particles of two or more pure substances which may be present in it in any ratio. Hence, their composition is variable. Pure substances forming mixture are called its components. Many of the substances present around you are mixtures. For example, sugar solution in water, air, tea, etc., are all mixtures.

Mixture (sub-divided)

  1. Homogeneous
  2. Heterogeneous.


In a homogeneous mixture, the components completely mix with each other. This means particles of components of the mixture are uniformly distributed throughout the bulk of the mixture and its composition is uniform throughout.
Sugar solution and air are examples of homogeneous mixtures.


The composition is not uniform throughout and sometimes different components are visible. For example, mixtures of salt and sugar, grains and pulses along with some dirt (often stone pieces), are heterogeneous mixtures.
πŸ‘‰ The components of a mixture can be separated by using physical methods, such as simple hand-picking, filtration, crystallisation, distillation, etc

Pure substances

Constituent particles of pure substances have fixed composition. Copper, silver, gold, water and glucose are some examples of pure substances. Glucose contains carbon, hydrogen and oxygen in a fixed ratio and its particles are of same composition. 
πŸ‘‰ Its constituents—carbon, hydrogen, and oxygen—cannot be separated by simple physical methods.

Pure substances (Sub-divided)

  1. Elements
  2. Compounds


Particles of an element consist of only one type of atoms. These particles may exist as atoms or molecules.
Some elements, such as sodium or copper, contain atoms as their constituent particles, whereas, in some others, the constituent particles are molecules which are formed by two or more atoms. For example, hydrogen, nitrogen and oxygen gases consist of molecules, in which two atoms combine to give their respective molecules. 
A representation of atoms and molecules
 A representation of atoms and molecules (Image credit: NCERT)


When two or more atoms of different elements combine together in a definite ratio, the molecule of a compound is obtained.


The combination of elements to form compounds is governed by the following five basic laws
  1. Law of Conservation of Mass
  2. Law of Definite Proportions
  3. Law of Multiple Proportions
  4. Gay Lussac’s Law of Gaseous Volumes
  5. Avogadro’s Law

1. Law of Conservation of Mass

In all physical and chemical changes, there is no net change in mass during the process. Hence, he reached the conclusion that matter can neither be created nor destroyed. 
In fact, this was the result of the exact measurement of masses of reactants and products, and carefully planned experiments performed by Lavoisier

2. Law of Definite Proportions

 A given compound always contains exactly the same proportion of elements by weight.

3. Law of Multiple Proportions

If two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element, are in the ratio of small whole numbers.
For example, hydrogen combines with oxygen to form two compounds, namely, water and hydrogen peroxide.

Here, the masses of oxygen (i.e., 16 g and 32 g), which combine with a fixed mass of hydrogen (2g) bear a simple ratio, i.e., 16:32 or 1: 2

4. Gay Lussac’s Law of Gaseous Volumes

When gases combine or are produced in a chemical reaction they do so in a simple ratio by volume, provided all gases are at the same temperature and pressure.
Thus, 100 mL of hydrogen combine with 50 mL of oxygen to give 100 mL of water vapour.

Thus, the volumes of hydrogen and oxygen which combine (i.e., 100 mL and 50 mL) bear a simple ratio of 2:1.

5. Avogadro’s Law

Equal volumes of all gases at the same temperature and pressure should contain equal number of molecules. 

Two volumes of hydrogen react with one volume of oxygen to give two volumes of water vapour
(Image credit: NCERT))

Atomic Masses

In 1961 for a universally accepted atomic mass unit, carbon-12 isotope was chosen as the standard reference for measuring atomic masses. One atomic mass unit is a mass unit equal to exactly one-twelfth (1/12th) the mass of one atom of carbon-12. The relative atomic masses of all elements have been found with respect to an atom of carbon-12.
 One atomic mass unit is defined as a mass exactly equal to one-twelfth of the mass of one carbon - 12 atom.

1 amu `= 1.66056×10^{–24}` g

Mass of an atom of hydrogen `= 1.6736×10^{–24}` g

At present, ‘amu’ has been replaced by ‘u’, which is known as unified mass.

Average Atomic Mass

Many naturally occurring elements exist as more than one isotope.  For example, carbon has the following three isotopes with relative abundances and masses as shown against each of them.

Image credit: NCERT

From the above data, the average atomic mass of carbon will come out to be: 
`(0.98892) (12 u) + (0.01108) (13.00335 u) + (2 × 10^{–12}) (14.00317 u) = 12.011 u`

Molecular Mass

Molecular mass is the sum of the atomic masses of the elements present in a molecule. It is obtained by multiplying the atomic mass of each element by the number of its atoms and adding them together.
Molecular mass of methane, 
`CH_4 = (12.011 u) + 4 (1.008 u) = 16.043 u`

Formula Mass

Some substances, such as sodium chloride, do not contain discrete molecules as their constituent units. In such compounds, positive (sodium ion) and negative (chloride ion) entities are arranged in a three-dimensional structure, as shown in Fig.
Image credit: NCERT

The formula, such as `NaCl`, is used to calculate the formula mass instead of molecular mass as in the solid state sodium chloride does not exist as a single entity. 
Thus, the formula mass of sodium chloride is the atomic mass of sodium `+` atomic mass of chlorine 
`= 23.0 u + 35.5 u = 58.5 u`


The number of particles (atoms, molecules or ions)  present in 1 mole of any substance is fixed, with a value of `6.022 × 10^{23}`. This is an experimentally obtained value. This number is called the Avogadro Constant or Avogadro Number (represented by `N_A`), named in honour of the Italian scientist, Amedeo Avogadro. 
1 mole ` = 6.022 × 10^{23}` in number,   
as, 1 dozen = 12 nos. 
1 gross = 144 nos.
The mass of 1 mole of a substance is equal to its relative atomic or molecular mass in grams. We have to take the same numerical value but change the units from ‘u’ to ‘g’. Molar mass of atoms is also known as gram atomic mass.
 For example, atomic mass of hydrogen `=1u`. So, gram atomic mass of hydrogen `= 1 g`. `1 u` hydrogen has only 1 atom of hydrogen 1 g hydrogen has 1 mole atoms, that is, `6.022 × 10^{23}` atoms of hydrogen.
18 u water has only 1 molecule of water, 18 g water has 1 mole molecules of water, that is, `6.022 × 10^{23}` molecules of water. 

1 mole `= 6.022 × 10^{23}` 
`=` Relative mass in grams.

The word “mole” was introduced around 1896 by Wilhelm Ostwald who derived the term from the Latin word moles meaning a ‘heap’ or ‘pile’. A substance may be considered as a heap of atoms or molecules.


`n =` No. of moles
`m =` given mass
`M =` Molar mass
`n = m/M`

`N =` Given no of particles
`N_A =` Avogadro number
`n = N/N_A`


An empirical formula represents the simplest whole-number ratio of various atoms present in a compound, whereas, the molecular formula shows the exact number of different types of atoms present in a molecule of a compound.
Empirical formula πŸ‘‰ `CH_2O`
Molecular formula πŸ‘‰ `C_6H_12O_6`
πŸ‘‰ If the mass per cent of various elements present in a compound is known, its empirical formula can be determined. The molecular formula can further be obtained if the molar mass is known. 

Q) A compound contains 4.07% hydrogen, 24.27% carbon, and 71.65% chlorine. Its molar mass is 98.96 g. What are its empirical and molecular formulas?
πŸ‘‰ Conversion of mass per cent to grams
Let 100 g of the compound
then, 4.07g hydrogen, 24.27g carbon and 71.65g chlorine are present

πŸ‘‰ Convert into the number of moles of each element.
Divide the masses by the respective atomic masses of various elements.
Moles of hydrogen `= frac{4.07 g}{1.008g} = 4.04`
Moles of carbon `= frac{24.27 g}{ 12. 01g} =  2. 021`
Moles of chlorine `= frac{71.65g}{35. 453g} =2. 021`

πŸ‘‰ Divide each of the mole values obtained above by the smallest number amongst them
Since `2.021` is smallest value, division by it gives a ratio of `2:1:1` for `H:C:Cl `.
`CH_2Cl` is, thus, the empirical formula of the above compound.

πŸ‘‰ Writing empirical formula
For `CH_2Cl`, empirical formula mass is 
`12.01 + (2 × 1.008) + 35.453 = 49.48 g`

 Divide Molar mass by empirical formula mass
Molar mass/ Empirical formula mass `= frac{98.96g}{49.48g} = 2 = n`
Multiply empirical formula by `n` obtained above to get the molecular formula
Empirical formula `= CH_2Cl, n = 2`. Hence molecular formula is `C_2H_4Cl_2`.

Reactions in Solutions

The concentration of a solution or the amount of substance present in its given volume can be expressed in any of the following ways.

1. Mass per cent or weight per cent (w/w %) 
2. Mole fraction 
3. Molarity 
4. Molality

1. Mass per cent 

It is obtained by using the following relation:

2. Mole Fraction

It is the ratio of number of moles of a particular component to the total number of moles of the solution. If a substance ‘A’ dissolves in substance ‘B’ and their number of moles are `n_A` and `n_B`, respectively, then the mole fractions of `A` and `B` are given as:

3. Molarity

It is the most widely used unit and is denoted by M. It is defined as the number of moles of the solute in 1 litre of the solution. Thus,

4. Molality

It is defined as the number of moles of solute present in 1 kg of solvent. It is denoted by m.

  1. NCERT Chemistry Class 11


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