Some Basics Concepts of Chemistry

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Chapter 1 : Some Basics Concepts of Chemistry

1. Introduction arrow_upward

  • Chemistry is the branch of science that studies the composition, properties and interaction of matter.

  • 1.1 Importance of Chemistry

  • Chemistry plays an important role in daily life.
  • Chemical principles are important in diverse areas, such as: weather patterns, functioning of brain and operation of a computer.
  • Chemical industries manufacturing fertilizers, alkalis, acids, salts, dyes, polymers, drugs, soaps, detergents, metals, alloys and other inorganic and organic chemicals, including new materials, contribute in a big way to the national economy.

  • 1.2 Nature of Matter arrow_upward


  • Anything which has mass and occupies space is called matter.
  • Matter can exist in three physical states:
  • Solid:
  • Solids have definite volume and definite shape.
  • Liquid:
  • Liquids have definite volume but not the definite shape. They take the shape of the container in which they are placed.
  • Gas:
  • Gases have neither definite volume nor definite shape. They completely occupy the container in which they are placed.

  • 1.3 Properties of Matter and their Measurement arrow_upward

  • Every substance has unique or characteristic properties. These properties can be classified into two categories – physical properties and chemical properties.
  • Physical Properties:
  • Physical properties are those properties which can be measured or observed without changing the identity or the composition of the substance.
  • Some examples of physical properties are color, odour, melting point, boiling point, density etc.
  • Chemical Properties:
  • The measurement or observation of chemical properties requires a chemical change to occur.
  • The examples of chemical properties are characteristic reactions of different substances; these include acidity or basicity, combustibility etc.

  • 1.3.1 The International System of Units (SI)

  • The SI system has seven base units listed in the table given below:

  • Base Physical Quantity

    Symbol for Quantity

    Name of SI Unit

    Symbol for SI Unit













    Electric Current




    Thermodynamic temperature




    Amount of substance




    Luminous Intensity




    1.3.2 Mass and Weight:

  • Mass of a substance is the amount of matter present in it.
  • Weight is the force exerted by gravity on an object.
  • The mass of a substance is constant whereas its weight may vary from one place to another due to change in gravity.
  • The SI unit of mass is kilogram.
  • 1 kg = 1000 grams

  • Volume has the units of (length)3 . So in SI system, volume has units of m3 .
  • A common unit, liter (L) which is not an SI unit, is used for measurement of volume of liquids.
  • 1 L = 1000 mL

  • Density of a substance is its amount of mass per unit volume.
  • Temperature:
  • There are three common scales to measure temperature:
    • °C (degree Celsius)
    • °F (degree Fahrenheit)
    • K (Kelvin)
  • Here, K is the SI unit.

  • 1.4 Uncertainty in Measurement arrow_upward

  • Measurement uncertainty is a non-negative parameter characterizing the dispersion of the values attributed to a measured quantity. The uncertainty has a probabilistic basis and reflects incomplete knowledge of the quantity.

  • 1.4.1 Scientific Notation

  • A chemist has to deal with numbers as large as 602, 200, 000, 000, 000, 000, 000, and 000 for the molecules of 2 g of hydrogen gas or as small as 0.00000000000000000000000166 g mass of a H atom.
  • It offers a real challenge to do simple mathematical operations of addition, subtraction, multiplication or division with such numbers.
  • This problem is solved by using scientific notation for such numbers, i.e., exponential notation in which any number can be represented in the form N × 10n where n is an exponent having positive or negative values and N is a number (called digit term) which varies between 1.000... and 9.999....
  • Thus, 0.00016 can be written as 1.6 × 10–4 .

  • 1.4.2 Significant Figures

  • Every experimental measurement has some amount of uncertainty associated with it. However, one would always like the results to be precise and accurate.
  • Precision refers to the closeness of various measurements for the same quantity.
  • Accuracy is the agreement of a particular value to the true value of the result.
  • Consider the following example:
    • The true value for a result is 2.00 g and Student A, B and C takes the following measurement as shown in the table given below:





    Student A




    Student B




    Student C




    • The observations taken by Student A and B are neither precise nor accurate whereas the observation taken by Student C is both precise and accurate
  • The uncertainty in the experimental or the calculated values is indicated by mentioning the number of significant figures.
  • Significant figures are meaningful digits which are known with certainty. The uncertainty is indicated by writing the certain digits and the last uncertain digit.
  • Thus, if we write a result as 11.2 mL, we say the 11 is certain and 2 is uncertain and the uncertainty would be ±1 in the last digit.

  • Rules

  • There are certain rules for determining the number of significant figures. These are stated below:
    • All non-zero digits are significant. For example in 0.35 there are two significant figures.
    • Zeros proceeding to first non-zero digit are not significant. For example, 0.03 has one significant figure.
    • Zeros between two non-zero digits are significant.
    • Zeros at the end or right of a number are significant provided they are on the right side of the decimal point.
    • Counting numbers of objects, for example, 2 balls or 20 eggs have infinite significant figures.

    1.5 Laws of Chemical Combinations arrow_upward

    1.5.1 Law of Conservation of Mass

  • It states that matter can neither be created nor destroyed.

  • 1.5.2 Law of Definite Proportions

  • This law was given by, a French chemist, Joseph Proust. He stated that “a given compound always contains exactly the same proportion of elements by weight.”

  • 1.5.3 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”.

  • 1.5.4 Gay Lussac’s Law of Gaseous Volumes

  • This law was given by Gay Lussac in 1808. He observed that “when gases combine or are produced in a chemical reaction they do so in a simple ratio by volume provided all gases are at same temperature and pressure”.

  • 1.5.5 Avogadro Law

  • Avogadro proposed that “equal volumes of gases at the same temperature and pressure should contain equal number of molecules”.

  • 1.6 Dalton’s Atomic Theory arrow_upward

  • Dalton proposed the following:
    • 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.

    1.7 Atomic and Molecular Masses arrow_upward

    1.7.1 Atomic Mass

  • The present system of atomic masses is based on carbon - 12 as the standard. Here, Carbon - 12 is one of the isotopes of carbon.
  • One atomic mass unit is defined as a mass exactly equal to one-twelfth the mass of one carbon - 12 atoms.
  •     1 amu = 1.66056 × 10-24 g

        Mass of an atom of hydrogen =

    1.6736 × 10-24 g

  • In terms of amu, the mass of hydrogen atom is given by:
  • = 1.0078 amu

    = 1.0080 amu

    1.7.2 Average Atomic Mass

  • Many naturally occurring elements exist as more than one isotope. When we take into account the existence of these isotopes and their relative abundance (per cent occurrence), the average atomic mass of that element can be computed. For example, carbon has the following three isotopes with relative abundances and masses as shown against each of them.

  • Isotope

    Relative Abundance (%)

    Atomic Mass (amu)

    12 C



    13 C



    14 C

    2 × 10-10


  • 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

    1.7.3 Molecular 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.
  • Example:
  • Molecular mass of methane:
  • (CH4 ) = (12.011 u) + 4(1.008 u)

    = 16.043 u

    1.7.4 Formula Mass

  • Some substances such as sodium chloride do not contain discrete molecules as their constituent units.
  • 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, formula mass of sodium chloride = atomic mass of sodium + atomic mass of chlorine
  • = 23.0 u + 35.5 u = 58.5 u

    1.8 Mole Concept and Molar Masses arrow_upward

    Mole concept
  • One mole is the amount of a substance that contains as many particles or entities as there are atoms in exactly 12 g (or 0.012 kg) of the 12 C isotope.
  • The mole of a substance always contains the same number of entities, no matter what the substance may be.
    •     Mass of a Carbon-12 atom

    = 1.992648 × 10–23 g,

        One mole of carbon weighs 12 g,

    = 6.0221367 × 1023 atoms/mol

    Molar Mass
  • The mass of one mole of a substance in grams is called its molar mass. The molar mass in grams is numerically equal to atomic/molecular/formula mass in u.
  •     Molar mass of water = 18.02 g mol-1

    1.9 Percentage Composition arrow_upward

    Molar mass of water = 18.02 g

    = 11.18

    = 88.79

    1.9.1 Empirical Formula for Molecular Formula arrow_upward

    Empirical formula
  • An empirical formula represents the simplest whole number ratio of various atoms present in a compound.
  • Molecular formula
  • The molecular formula shows the exact number of different types of atoms present in a molecule of a compound.

  • 1.10 Stoichiometry and Stoichiometric Calculations arrow_upward

  • Stoichiometry, thus, deals with the calculation of masses (sometimes volumes also) of the reactants and the products involved in a chemical reaction.
  • Here, methane and dioxygen are called reactants and carbon dioxide and water are called products.
  • The coefficients 2 for O2 and H2 O are called stoichiometric coefficients.
  • Similarly the coefficient for CH4 and CO2 is one in each case. They represent the number of molecules (and moles as well) taking part in the reaction or formed in the reaction.

  • 1.10.1 Limiting Reagent arrow_upward

  • Many a time, the reactions are carried out when the reactants are not present in the amounts as required by a balanced chemical reaction.
  • In such situations, one reactant is in excess over the other. The reactant which is present in the lesser amount gets consumed after sometime and after that no further reaction takes place whatever be the amount of the other reactant present.

  • 1.10.2 Reactions in Solutions arrow_upward

  • The concentration of a solution or the amount of substance present in its given volume can be expressed in any of the following ways:
    • Mass per cent or weight per cent (w/w %)
    • Mole fraction
    • Molarity
    • Molality
    Mass per cent
  • It is obtained by following formula:
  • 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 nA and nB respectively; then the mole fractions of A and B are given as:
  • Molarity
  • It is defined as the number of moles of the solute in 1 litre of the solution.
  • Molality
  • It is defined as the number of moles of solute present in 1 kg of solvent. It is denoted by m.

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