CBSE Class 11 Chemistry Chapter-1

1.  Role of chemistry in different fields of life;
 Characteristics of three states of matter;
 Classify different substances into elements, compounds and mixtures;
 Define SI base units and list some commonly used prefixes;
 Precision and accuracy;
 Scientific notations and significant figures;
 Various laws of chemical combination;
 Atomic mass, molecular mass and formula mass;
 Mole concept;
 Calculate the mass per cent of different elements constituting a compound;
 Determine empirical formula and molecular formula for a compound;
 Stoichiometric calculations; limiting reagent, excess reagent;
 Concentration terms like molarity, molality, mole fraction etc.

2. Chemistry deals with the
composition, structure and
properties of matter.

3. Chemistry is an interdisciplinary subject :-

4. Matter :-
“Anything which occupies space and have mass is called matter.”
Three states of matter :-
 Solid
 Liquid
 Gas
 Plasma in stars
 BEC

5. Characteristics of matter :-

6. Characteristics of matter continued …

7. At the macroscopic or bulk level, matter can be classified as mixtures
or pure substances. These can be further sub-divided as

8. Mixture:- A mixture contains two or more substances present in it (in
any ratio) which are called its components.
 A homogeneous mixture, the components completely mix with each other
and its composition is uniform throughout.
 In heterogeneous mixtures, the composition is not uniform throughout.
 Pure substances have fixed composition.
 Pure substances can be further classified into elements and compounds.

9. Properties of matter :-
Physical properties :-
are those properties which
can be measured or
observed without changing
the identity or the
composition of the
substance. E.g. colour,
odour, melting point, boiling
point, density etc.
Chemical properties :-
are characteristic reactions of
different substances; these
include acidity or basicity,
combustibility etc.
The International System of Units (SI) :-
The International System of Units (in French Le Systeme
International d‘Unités -abbreviated as SI) was established by the 11th
General Conference on Weights and Measures (CGPM from Conference
Generale des Poids at Measures). The CGPM is an inter governmental treaty
organization created by a diplomatic treaty known as Metre Convention
which was signed in Paris in 1875.

10. The SI system has seven base units

11. SI Prefixes:-
 Prefix is chosen depending on the quantity to be measured.
 Conversion is through.
1. Adding “0s” to the power of 10.
e.g. : Distance from Delhi to Agra = 210 km
= 210000 m
2. Moving decimal points to the power of 10.
e.g. : Distance between dots = 2 cm = 0.02 m.
Quantities more than 1 Quantities less than 1

12. Accuracy :- How closely a measurement match the correct or
expected value.
Precision :- How closely repeated measurements match each other.
Significant figures are meaningful
digits which are known with
certainty.

13. Rules for Identification of significant figures :-
 Non zero digits are significant.
e.g. 245, 3 significant figures.
 Zeroes in the middle of a number are significant.
e.g. 101, 3 significant figures.
 Zeroes to the right of a digit after a decimal point are significant.
e.g. 0.1200, 4 significant figures.
 Zeroes between decimal point and a number are not significant.
e.g. 0.00150, 3 significant figures.
 Upon addition and subtraction with significant figures, the result
cannot have more figures to the right of the decimal point than
either of the original numbers.
 In multiplication and division with significant numbers, the answer
cannot have more significant figures than any of the original
numbers. e.g. A person weighing 75.3 kg gains 21% of his weight
in a month. What will be his total weight gain?

14.  Law of conservation of mass
 Law of definite proportion
 Law of multiple proportion
 Law of combining volumes (Gay Lussac’s Law)
 Avogadro's Law
LAWS OF CHEMICAL COMBINATIONS

15. Law of Conservation of Mass :-
“Matter can neither be created nor destroyed”.
R1 R2 P1 P2
m1 m2 m3 m4
Total mass of Total mass of
Reactants = m1 + m2 products = m3 + m4
According to law of conservation of mass :-
Total mass of reactant = Total mass of products
Therefore, m1 + m2 = m3 + m4

16. Law of definite proportions :-
“A chemical compound always contains
exactly the same proportion of elements
by weight”
2 g 16 g 18 g
By taking ratio 2:16 = 1:8
2H O H2O
Copper carbonate % OF
COPPER
% OF
OXYGEN
% OF
CARBON
Natural Sample 51.35 9.74 38.91
Synthetic Sample 51.35 9.74 38.91

17. Law of multiple proportions :-
“If two elements can combine to form more
than one compound, the mass of one
element that combines with the fixed mass
of the other element is in the ratio of small
whole numbers.”

18. 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 the gases are
at same temperature and pressure.”
3 Vol. 1 Vol. 2 Vol.
Ratio = 3:1:2
H2
N2 NH3

19. Avogadro law :-
“At the same temperature and pressure,
equal volumes of gases contain equal
number of molecules.”
H2
N2 NH3
3 Vol. 1 Vol. 2 Vol.
3n molecules n molecules 2n molecules
of Hydrogen of Nitrogen of Ammonia
3 molecules 1 molecules 2 molecules
Of Hydrogen of nitrogen of ammonia

20. DALTON’S ATOMIC THEORY :-
 Matter consists of indivisible atoms.
 All the atoms of a given element have
identical properties including identical
mass. Atoms of different elements differ
in mass.
 Compounds are formed when atoms of
different elements combine in a fixed
ratio.
 Chemical reactions involve reorganization
of atoms. These are neither created nor
destroyed in a chemical reaction.

21. ATOMIC AND MOLECULAR MASSES :-
Atomic Mass :-
 The atomic mass is a weighted average of all of the
isotopes of that element, in which the mass of each
isotope is multiplied by the abundance of that particular
isotope.
 One atomic mass unit is defined as a mass exactly
equal to one twelfth the mass of one carbon – 12 atom.
1 amu = 1.66056 x 10-24 g
Today, ‘amu’ has been replaced
by ‘u’ which is known as unified
mass.

22. Molecular Mass & Formula Mass :-
 Molecular mass is the sum of 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.
 Formula mass is defined as the sum of
atomic masses of the ions present in the formula unit
of an ionic compound.
e.g. NaCl.

23. Mole Concept :-
 A mole of a substance or a mole of particles is defined as exactly
6.02214076 ×10²³ particles, which may be atoms, molecules, ions,
or electrons.
 In short, for particles 1 mol = 6.02214076 ×10²³.

24. PERCENTAGE COMPOSITION :-
Percent composition is the percentage by mass of each element in a
compound.
 Find the molar mass of all the elements in the compound in grams
per mole.
 Find the molecular mass of the entire compound.
 Divide the component's molar mass by the entire molecular mass.
 You will now have a number between 0 and 1. Multiply it by 100%
to get percent composition.
e.g.
 What is the percentage of carbon, hydrogen and oxygen in
ethanol?
 What is the percent composition by mass of aspartame
(C14H18N2O5), the artificial sweetener NutraSweet?

25. Empirical Formula for Molecular Formula
 Conversion of mass per cent
to grams.
 Convert into number moles
of each element.
 Divide the mole value
obtained above by the
smallest number to get a
simple ratio.
 Write empirical formula by
mentioning the numbers
after writing the symbols of
respective elements.
 e.g. benzene and acetylene.
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.

26. STOICHIOMETRY AND STOICHIOMETRIC CALCULATIONS
 In a chemical reaction, it is very
important to note the amount of
reactants required to prepare a
particular amount of product.
 Such a calculation is called
stiochiometry.
Limiting reactant :-
 The reactant that is completely
consumed in a chemical reaction
is called limiting reactant.
 e.g.
 Calculate the amount of water (g) produced by the combustion of
16 g of methane
 How many moles of KCl will be produced from 15.0 g of KClO3?
 How many moles of AgCl will be produced from 60.0 g of AgNO3​,
assuming NaCl is available in excess?

27. Reactions in Solutions :-
 Mass Percent is the mass of the
solute in grams per 100 grams of the
solution.
 Volume percent is the number of
units of volume of the solute per 100
units of the volume of solution.
 Mole Fraction is the ratio of number
of moles of a particular component to
the total number of moles of the
solution.
 Molarity 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.
 Molality is defined as the number of
moles of solute present in 1 kg of
solvent. It is denoted by m.

28. Molarity vs Molality :-
Molarity is temperature dependant, it changes with change in volume.
 Relation between Molarity, Molality and
Mole fraction
e.g.
 How many grams of NaOH are there in
500 mL of a 2 M NaOH solution?
 What is the concentration of a solution
in percentage (w/v) made by dissolving
4.6 g NaCl in water to have 250 mL of
salt solution?
 Calculate the mole fraction, molarity
and molality of 0.26 L of acetic acid
solution. The solution is composed of
14.1 g of acetic acid and 250 g of water.
The molecular mass of acetic acid is
60.
The Dilution Equation
Where,