Some Basic Concepts Of Chemistry Class 11 Notes Chemistry

some basic concepts of chemistry class 11 notes

This article is on the Some Basic Concepts Of Chemistry Class 11 Notes of Chemistry. The notes on Some Basic Concepts Of Chemistry of class 11 have been prepared with great care keeping in mind the easiness of it for the students. This article provides the revision notes of the Some Basic Concepts Of Chemistry chapter of Class 11 Chemistry for the students so that they can give a quick glance of the chapter.

Some Basic Concepts Of Chemistry

Chemistry is defined as that branch of science which deals with the study of composition, structure and properties of matter and the transformations which the matter undergoes under different conditions and the laws which govern these changes.

A matter is defined as anything that occupies space, possesses mass and the presence of which can be felt by any one or more of our five senses. A matter is classified into three types: Solid, liquid and gas. A solid is composed of matter where the particles have strong attractive forces, are close together and have fixed positions. Solids have a definite shape and a definite volume.  A liquid is composed of matter where the particles are close together but have weak attractive forces so that, they are able to move past one another therefore, liquids have a definite volume but no definite shape. A gas is made up of atoms or molecules with large interparticle distances, which move independently in all directions in random motion. Gases neither have definite volume nor definite shape.

arrangement of particles in solid, liquid and gaseous state

A matter is classified as pure substances and mixtures. Pure substances are either elements or compounds and are always homogeneous in composition. Mixture always contains two or more substances which may be either homogeneous or heterogeneous.

  1. A homogeneous mixture has a uniform composition, appearance old properties throughout as the components completely mix with each other.
  2. A heterogeneous mixture has two or more physically distinct phases and the composition is not uniform throughout.

classification of matter

Physical properties are those which can be measured or observed without changing the identity or composition of the substance, e.g., mass, melting point, boding point and density, etc. Many physical properties of the substance are quantitative in nature.

Chemical properties are those in which a chemical change in the substances occurs, e.g., acidity or basicity, combustibility, etc.

International System of Units

The International System (SI, In French Systeme Internationale) is a decimal system of a unit for measurement of mass, length, time (fundamental units) and other physical quantities (desired units).

base physical quantities and their units

Mass and Weight

The mass is the amount of matter present in a substance and is measured very accurately in the laboratory by using an analytical balance. The SI unit of mass is kilogram. The weight of a substance is the force exerted by gravity on an object.

A density of a substance is its amount of mass per unit volume. Its SI unit is kg.


The SI unit of temperature is K (Kelvin). It is also measured in 0C (degree Celsius) and 0F (degree Fahrenheit).

0F  =9/5(0C) + 32

K = °C + 273.15

Uncertainty in Measurement

Measurement is the quantity, dimensions or extent of something, usually in comparison to a specific unit. A unit is a fixed quantity chosen as the standard of measurement. All measurements show some amount of uncertainty and are indicated by the number of significant figures in the measurement. The accuracy of any such measurement depends upon (i) the accuracy of the measuring device used, and (ii) the skill of its operator.

Scientific Notation

Scientific notation is the exponential notation in which any number can be represented in the form N x 10n where it is an exponent having positive or negative values and N can vary between 1 to 10.

Significant figures

If the average value of different measurements is close to the correct value, the measurement is said to be accurate (the individual measurement may not be close to each other). If the values of different measurements are close to each other and hence close to their average value, the measurement is said to be precise (the average value of different measurement may not be close to the correct value).

The total number of digits in a number including the last digit whose value is uncertain is called the number of significant figures.

(1) Common rules for counting significant figures. Following rules are applied for counting the number of significant figures in a given quantity

Rule 1. All non zero digits, as well as the zeros between the non zero digits, are significant.

Example: 189 has three significant figures, 1007 has four significant figures, 1.0809 has five significant figures and 0.48 has two significant figures.

Rule 2. If a number ends in zeros but these zeros are to the right of the decimal point, then these zeros are significant.

Example: 5.0 has two significant figures, 1.50 has three significant figures, 5.500 has four significant figures and 0.0200 has three significant figures.

Rule 3. If a number ends in zeros but these zeros are not to the right of the decimal point, then these zeros may or may not be significant.

Example: 10500 may have three, four or five significant figures. This ambiguity is removed by expressing the value in an exponential form.

1.05 x 104, which has three significant figures, 1.050 x 104, which has four significant figures, 1.0500 x 104, which has five significant figures.

Rule 4. Zeros to the left of the first non zero digit in a number are not significant. (They simply indicate the position of the decimal point).

Example: 0.05 has only one significant figure, 0.0045 has two significant figures.

Thus, we can say that by changing the position of the decimal point, the number of significant digits in the results does not change.

(2) Rounding off: Following rules are applied while rounding off measurements:

Rule 1. If the digit just next to the last digit to be retained is less than 5, the last digit is taken as such and all other digits on its right are dropped.

Example: 7.82567 is rounded off to 7.8, 3.9434 is rounded off to 3.9.

Rule 2. If the digit is greater than 5, the last digit to be retained is increased by 1 and all other digits on its right are dropped.

Example: 6.87 is rounded off to 6.9, 12.78 is rounded off to 12.8.

Rule 3. If the digit is equal to 5, the last significant figure is left unchanged if it is even and is increased by 1 if it is odd.

Example: 3.250 becomes 3.2, 12.650 becomes 12.6, 3.750 is rounded off to 3.8 and 16.150 is rounded off to 16.2.

(3) The significant figure in a calculation

(i) Addition and subtraction:

(a) Every quantity should be changed into the same unit.
(b) If a quantity is expressed in the power of 10, then all the quantities should be changed into the power of 10.
(c) The result obtained after addition or subtraction, the number of figures should be equal to that of least, after the decimal point.

(ii) Multiplication and division

(a) The number of significant figures will be the same if any number is multiplied by a constant.
(b) The product or division of two significant figures will contain the significant figures equal to that of least.

Laws of chemical combination

The chemical reactions take place according to certain laws, called the Laws of chemical combination. These are:

(a) Law of conservation of mass: “In all physical and chemical changes, the total mass of the reactants is equal to that of the products” or “matter can neither be created nor destroyed.”

(b) Law of constant composition/definite proportion: “A chemical compound is always found to be made up of the same elements combined together in the same fixed ratio by weight”.

(c) Law of multiple proportions: “When two elements combine together to form two or more chemical compounds, then the weight of one of the elements which combine with a fixed weight of the other bear a simple ratio to one another”.

(d) Law of reciprocal proportions: The ratio of the weights of two elements A and B which combine with a fixed weight of the third element C is either the same or a simple multiple of the ratio of the weights of A and B which directly combine with each other.

(e) Gay-Lussac’s law of gaseous volumes: “When gases react together, they always do so in volumes which bear a simple ratio to one another and to the volumes of the products, if gaseous, all measurements are made under the same conditions of temperature and pressure”

(f) Avogadro Law: Equal volumes of all gases/ vapours under similar conditions of temperature and pressure contain an equal number of molecules.

Dalton’s Atomic Theory:

In 1808, John Dalton put forward a theory known as Dalton’s atomic theory. The main points of this theory are as follows-

 (i) A matter is made up of extremely small indivisible particles called atoms.

(ii) Atoms of the same element are identical in all respects, i.e., size, shape and mass.

(iii) Atoms of different elements have different masses, sizes and also possess different chemical properties.

(iv)  Atoms can neither created nor destroyed.

(v) Atoms of same or different elements combine together to form.

Atomic and Molecular masses

Atomic Mass:

The atomic mass of an element is defined as the average relative mass of its atoms as compared to the mass of a carbon atom taken as 12. Note that the atomic weight of an element is a relative weight of one atom and not the absolute weight. In other words, the atomic mass (or atomic weight) of an element is the average relative mass of its atoms on a scale in which an atom of carbon-12 has a mass of exactly 12 u (where u or amu is unified mass) and 1 u = 1.66056 x 10-24g.

Average atomic mass:

It is the sum of the products of fractional abundances of the isotopes and their corresponding mass numbers.

Average atomic mass = (m x a) + (m x b)/m + n

where ‘a’ and ‘are the atomic masses of any two isotopes with fractional abundances m and n respectively.

One atomic mass unit (amu) is equal to 1/12th of the mass of an atom of carbon-12 isotope.

Molecular mass and Formula mass:

The sum of atomic masses of the elements present in a molecule is called molecular mass and in an ionic compound is called formula mass.

Mole concept and Molar Masses

Some Basic Concepts Of Chemistry class 11 notes

A mole is that amount of the substance which contains as many elementary entities as there are atoms in exactly 12 g of 12C or a mole is defined as that amount of the substance which has a mass equal to gram atomic mass if the substance is atomic or gram molecular mass if the substance is molecular. This number is also known as Avogadro’s number (NA). Avogadro’s number = 6.023 × 1023 molecule/mole. Thus, 1 mole of an entity contains NA particles of that entity. The mass of one mole of a substance in grams is called its molar mass, which is numerically equal to atomic/ molecular/ formula mass expressed in u.

Percentage composition

Percentage composition of the compound is the relative mass of each of the constituent element in 100 parts of it. If the molecular mass of a compound is M and B is the mass of an element in the molecule, then

Percentage of element = mass of that element in the compound x 100/ molar mass of the compound.

Empirical Formula for Molecular Formula:

The empirical formula is the simplest whole number ratio of atoms present in a compound. The molecular formula is the chemical formula which represents the true formula of its molecule. It expresses the actual number of atoms of various elements present in a molecule of a compound. Two or more compounds may have the same empirical formula but the molecular formula cannot be the same.

Determination of empirical formula:

The empirical formula of a molecule is determined using the % of elements present in it. The following method is adopted.

Element                 % Relative no. of atoms or %/at. wt. =              Simplest Ratio            Empirical Formula

Relative no. of atoms: Divide the percentage of each element present in the compound by its atomic weight. This gives the relative no. of atoms of the element in the molecule.

Simplest ratio: Find out the lowest value of relative no. of atoms and divide each value of relative no. of atoms by this value to estimate the simplest ratio of elements.

Empirical formula: Write all constituent atoms with their respective no. of atoms derived in simplest ratio. This gives the empirical formula of the compound.

Molecular formula: Molecular formula = n´ empirical formula where ‘ n‘ is the whole no. obtained by

             n = molecular weight of compound/empirical formula weight of the compound


The quantitative aspect, dealing with mass and volume relations between reactants and products is termed stoichiometry. Consider, for example, the reaction represented by a balanced chemical equation:

Chemical Equation:    2 H2 (g)          +         O2 (g)            →           2 H2O

Mole ratio:                    2 mol or                    1 mol or                       2 mol or

Molecule ratio :    2 x  6.023 x 1023         1 x  6.023 x 1023            2 x  6.023 x 1023
molecules                     molecules                       molecules

                                 or 2molecules             or 1molecules                  or 2molecules

Weight ratio:                 4g                               32g                                 36g

Volume ratio:             2 vol                            1 vol                              2 vol (valid only for the gaseous state at same P and T)

The given reaction suggests the combination ratio of reactants and the formation ratio of products in terms of:

(a) Mole ratio: 2 mol H2 reacts with 1 mol of O2 to form 2 mol of H2O vapours.

(b) Molecular ratio: 2 molecule of H2 reacts with 1 molecule of O2 to form 2 molecules of H2O vapours.

(c) Weight ratio: 4 g H2 reacts with 32 g O2 to form 36 g of H2O vapours.

(d) Volume ratio: In gaseous state 2 volume H2 reacts with 1 volume O2 to form 2 volume H2O vapours at same conditions of P and T. Therefore, coefficients in the balanced chemical reaction can be interpreted as the relative number of moles, molecules or volume (if reactants are gases) involved in the reaction. These coefficients are called stoichiometrically equivalent quantities and may be represented as:
2 mol H2 ≡ 1 mol O2 ≡ 2 mol H2O
Or Mole of H2: Mole of O2: Mole of H2O = 2: 1: 2

Where the symbol ≡ is taken to mean ‘stoichiometrically equivalent to’.

The Limiting Reagent

The reagent producing the least number of moles of products is the limiting reagent. For example, consider a chemical reaction given below, containing 10 mol of H2 and 7 mol of O2. Since, 2 mol H2 reacts with 1 mol O2, thus,

2H2 (g) + O2 (g) —–>  2H2O(V)

Moles before reaction           10             7                     0

Moles after reaction               0              2                    10

It is thus, evident that the reaction stops only after consumption of 5 moles of O2 since no further amount of H2 is left to react with unreacted O2. The substance that is completely consumed in a reaction is called limiting reagent because it determines or limits, the amount of product. The other reactants present in excess are sometimes called as excess reagents.


  1. Mass Percentage or Percent by Mass:

%(w/w) Mass percentage of solute = Mass of solute x 100/Mass of solution

  1. Mole Fraction: It is the fractional part of the moles that are contributed by each component to the total number of moles that comprises the solution. In containing nA moles of solvent and nB moles of solute.

                                              Mole fraction of B     xB= nB/ nA + nB

                                              Mole fraction of A  xA= nA/ nA + nB

  1. Molarity: It is expressed as moles of solute contained in one litre of solution or it is also taken as millimoles
    of solute in 1000 cc(mL) of the solution. It is denoted by M.

                                      Molarity = Moles of solute/ Litres of solution

                                                        = Millimoles of solute/Millilitres of solution

  1. Molality: It is the number of moles present in the 1kg solvent.

            Molality(m) = No. of moles of solute/Weight (in kg) of solvent

Let wA grams of the solute of molecular mass mA be present in wB grams of the solvent, then

                                                          Molality(m) = wA x 1000/ mA  x wB

This article has highlighted some basic concepts of chemistry in the form of notes for class 11 students in order to understand the chapter with ease. The notes on some basic concepts of chemistry have not only been prepared for class 11 but also for the different competitive exams such as iit jee, neet, etc.

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