Some Basic Concepts Of Chemistry Class 11 Notes Chemistry Chapter 1 - CBSE

Chapter : 1

What Are Some Basic Concepts Of Chemistry ?

The dot mark field are mandatory, So please fill them in carefully
To download the complete Syllabus (PDF File), Please fill & submit the form below.


    Anything that occupies space and has mass is called Matter.

    Physical Nature Of Matter

    • Matter is made up of particles.
    • The particles of matter are very small they are small beyond our imagination.

    Characteristics Of Particles Of Matter

    • Particles of matter have space between them.
    • Particles of matter are continuously moving.
    • Particles of matter attract each other.

    States Of Matter

    • There are three states of matter namely solid, liquid and gas.
    • The main difference between these states of matter are given below:



    States Of Matter

    • Definite shape
    • Distinct boundaries and fixed volume
    • Not compressible
    • Cannot flow
    • Very less inter-particle space
    • Maximum density
    • Negligible rate of diffusion
    • Inter-particle attraction is maximum
    • No definite shape
    • Fixed volume
    • Can be compressible
    • Can flow
    • Inter-particle spaces are more than in solids
    • Density is between that of the solids and gases
    • Rate of diffusion depends on inter-particle attraction
    • Inter-particle attraction is maximum
    • No definite shape
    • No fixed volume
    • Highly compressible
    • Can flow
    • Large inter-particle space is available
    • Minimum density
    • Maximum rate of diffusion
    • Inter-particle attraction is minimum



    It contains particles of two or more pure substances which may be present in it in any ratio.

    Homogeneous Mixtures

    The components completely mix with each other

    Heterogeneous Mixtures

    The composition is not uniform throughout  and sometimes different components are visible.

    Pure Substances

    They have characteristics different from mixtures, i.e., constituent particles of pure substances have fixed composition.


    Particles of an element consist of only one type of atoms. These particles may exist as atoms or molecules.


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

    Properties Of Matter And Their Measurements

    • Physical Properties can be measured or observed without changing the identity or the composition of the
      substance. Some examples of physical properties are colour, odour, melting point, boiling point etc.
    • Chemical Properties require a chemical change to occur. The examples of chemical properties are
      characteristic reactions of different substances. These include acidity, basicity, combustibility etc.

    The International System Of Units (SI)

    The SI system has seven base units listed below. These units pertain to the seven fundamental scientific quantities. The other physical quantities, such as speed, volume, density, etc., can be derived from these

    Base Physical Quantities and their Units

    Base Physical Quantity Symbol for Quantity Name of SI Unit Symbol for SI Unit
    Length l metre m
    Mass m kilogram kg
    Time t second s
    Electric current I ampere A
    Thermodynamic temperature T kelvin K
    Amount of substance n mole mol
    Luminous intensity lv candela cd


    Mass of a substance is the amount of matter present in it, while weight is the force exerted by gravity on an object. The mass of a substance is constant. The mass of a substance can be determined accurately in the laboratory by using an analytical balance. SI unit of mass is kilogram.


    It is the force exerted by gravity on an object. Weight of substance may vary from one place to another due to change in gravity.


    Volume means the space occupied by matter. It has the units of (length)3 . In SI units, volume is expressed in metre3 (m3). However, a popular unit of measuring volume, particularly in liquids is litre (L) but it is not an S.I. unit. Mathematically, 1L = 1000 mL = 1000 cm3 = 1dm3 . Volume of liquids can be measured by different devices like burette, pipette, cylinder, measuring flask etc.


    Density of a substance is its amount of mass per unit volume. Its SI unit is g cm–3.


    There are three scales in which temperature can be measured. These are known as Celsius scale (°C), Fahrenheit scale (°F) and Kelvin scale (K). Thermometres with Celsius scale are calibrated from 0°C to 100°C. Thermometres with Fahrenheit scale are calibrated from 32°F to 212°F Kelvin’ scale of temperature is S.I. scale and temperature on this scale is shown by the sign K. The temperature on two scales are related to

    each other by the relationship :

    $$\degree\text{F}=\frac{9}{5}\space(\degree\text{C}) + 32$$

    The Kelvin scale is related to Celsius scale as follows: K = °C + 273.15

    Significant Figures

    Significant figures are meaningful digits which are known with certainty. There are certain rules for determining the number of significant figures. These are stated below :

    • All non-zero digits are significant. For example, in 285 cm, there are three significant figures and in 0.25 mL,
      there are two significant figures.
    • Zeros preceding to first non-zero digit are not significant. Such zeros indicate the position of decimal point.
      For example, 0.03 has one significant figure and 0.0052 has two significant figures.
    • Zeros between two non-zero digits are significant. Thus, 2.005 has four significant figures.
    • Zeros at the end or right of a number are significant provided they are on the right side of the decimal point.
      For example, 0.200 g has three significant figures.
    • Counting numbers of objects. For example, 2 balls or 20 eggs have infinite significant figures as these are
      exact numbers and can be represented by writing infinite number of zeros after placing a decimal. i.e., 2 = 2.000000 or 20 = 20.000000

    Laws Of Chemical Combination

    Law of Conservation of Mass

    The law was established by a French chemist, A. Lavoisier. The law states: In all physical and chemical changes, the total mass of the reactants is equal to that of the products, i.e., matter can neither be created nor destroyed.

    Law of Definite Proportions

    This law was given by, a French chemist, Joseph Proust. The law states that, a given compound always contains exactly the same proportion of elements by weight.

    Law of Multiple Proportions

    This law was proposed by Dalton in 1803. According to this law, 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.

    Gay Lussac’s Law of Gaseous Volumes

    This law was given by Gay Lussac in 1808. The law states that, under similar conditions of temperature and pressure, whenever gases combine, they do so in volumes which bear simple whole number ratio with each other and also with the gaseous products.

    Avogadro’s Law

    This law was proposed by Avogadro in 1811. According to this, equal volumes of all gases at the same temperature and pressure should contain equal number of molecules.

    Dalton’s Atomic Theory

    In 1808, Dalton published ‘A New System of Chemical Philosophy’, in which he proposed the following :

    • Matter consists of indivisible atoms.
    • All 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 reorganisation of atoms. These are neither created nor destroyed in a chemical reaction.

    Mole Concept

    It is found that one gram atom of any element contains the same number of atoms and one gram molecule of any substance contains the same number of molecules. This number has been experimentally determined and found to be equal to 6.022 x 1023. The value is generally called Avogadro’s number or Avogadro’s constant. It is usually represented by NA :

    Avogadro’s Number, NA= 6.022 × 1023

    Percentage Composition

    $$\text{Mass} \%\space \text{of an element =}\\\frac{\text{mass of that element in the compound × 100}}{\text{molar mass of the compound}}$$

    Empirical Formula

    The formula of the compound which gives the simplest whole number ratio of the atoms of various elements present in one molecule of the compound. For example, the formula of hydrogen peroxide is H202 . In order to express its empirical formula, we have to take out a common factor 2. The simplest whole number ratio of the atoms is 1 : 1 and the empirical formula is HO.

    Molecular Formula

    The formula of a compound which gives the actual ratio of the atoms of various elements present in one molecule of the compound.

    Molecular formula = n x Empirical formula

    Where n is the common factor and also called multiplying factor. The value of n may be 1, 2, 3, 4, 5, 6 etc.

    In case n is 1, Molecular formula of a compound = Empirical formula of the compound.

    Limiting Reactant/Reagent

    In a chemical equation, the reactants present are not the amount as required according to the balanced equation. The amount of products formed then depends upon the reactant which has reacted completely. This reactant which reacts completely in the reaction is called the limiting reactant or limiting reagent. The reactant which is not consumed completely in the reaction is called excess reactant.

    Reactions In Solutions

    When the reactions are carried out in solutions, the amount of substance present in its given volume can be expressed in any of the following ways:

    Mass Percent

    It is obtained by using the following relation :

    • Weight percent (% w/W) = (Weight of Solute / Weight of Solution) × 100
    • Volume percent (% V/V) = (Volume of Solute / Volume of Solution) × 100

    Mole fraction

    It is the ratio of number of moles of a particular component to the total number of moles of the solution.

    For n moles of solute and N moles of solvent, the mole fraction is,

    $$\text{Mole fraction of solute}(X_{\text{solute}})\\=\frac{n}{\text{n + N}}\\\text{Mole fraction of solvent (X}_\text{solvent})\\=\frac{n}{\text{n + N}}$$

    Xsolute + Xsolvent = 1


    It is defined as the number of moles of solute in 1 litre of the solution.

    $$\text{Molarity (M) = }\\\frac{\text{Number of moles of solute}}{\text{Volume of Solutions in litres}}$$


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

    $$\text{Molarity (M) =}\\\frac{\text{Number of moles of solute}}{\text{Mass of solvent in kg}}$$