Some Basic Concepts of Chemistry Class 11 Notes Chapter 1

Introduction, Matter, Element, Compound, Mixture, Atom, Dalton's Atomic Theory, Atomic Weight and Molecular Weight, The Mole, Chemical Formulae


In this chapter, we shall study about Some Basic Concept of Chemistry like matter, element, moles, molecules and chemical properties of matter.

Chemistry is the science of molecules and their transformations which deals with the study of matter, its composition, the changes that matter undergoes and the relation between changes in composition and changes in energy. Chemistry plays an important role in meeting human needs for food, health care products.

Importance of Chemistry

Chemistry plays a central role in science and is often intertwined with other branches of science.

Principles of chemistry are applicable in diverse areas, such as weather patterns, functioning of brain and operation of a computer, production in chemical industries, manufacturing fertilisers, alkalis, acids, salts, dyes, polymers, drugs, soaps, detergents, metals, alloys, etc., including new material.

Chemistry contributes in a big way to the national economy. It also plays an important role in meeting human needs for food, healthcare products and other material aimed at improving the quality of life.


Matter is any thing that occupies space, has mass, offer resistance and can be perceived of directly by our senses. For example, book, pen, pencil, water, air, all living beings, etc.

Some Basic Concepts of Chemistry


Element is the purest form of matter. It is made up of only one type of atoms, ex.- carbon, iron, copper, oxygen etc.


Compound is the substance which is made up of two or more elements combined together in a fixed ratio by their weight e.g., carbon dioxide.

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Mixture is the substance which is made up of two or more substances in any ratio. e.g., Sugar + Water, Sodium Chloride + Water, Sand + Water

On the basis of composition, mixtures are of following type :

1. Homogeneous mixture : The mixture which has uniform composition through out e.g., sugar solution.

2. Heterogeneous mixture : The mixtures which do not have uniform composition through out. e.g. sand in water.


Atom is the smallest particle which may or may not exist free but takes part in chemical reaction. Atom word means not to be cut. Ex- H, Na, O etc.

International System of Units (S.I.)

The international system of units (in French Le Systeme International d' Unités - abbreviated as SI) was established in 1960 by the 11th general conference on weights and measures. SI system is a modification of metric system and has seven base units pertaining to the seven fundamental scientific quantities.

Some Basic Concepts of Chemistry

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Definition of SI Base Units

(i). Metre : The metre is the length of the path travelled by light in vacuum during a time interval of 1/(299792458) of a second.

(ii). Kilogram : The kilogram is the unit of mass; it is equal to the mass of the international prototype of the kilogram.

(iii). Second : The second is the duration of 9192631770 periods of the radiation corresponding to the transition between the two hyperfine levels of the ground state of the caesium-133 atom.

(iv). Ampere : The ampere is that constant current which, if maintained in two straight parallel conductors of infinite length, of negligible circular cross-section, and placed 1 metre apart in vacuum, would produce between these conductors a force equal to 2 × 10-7 newton per metre of length.

(v). Kelvin : The kelvin, unit of thermodynamic temperature, is the fraction 1/273.16 of the thermodynamic temperature of the triple point of water.

(vi). Mole : The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12; its symbol is “mol”.

(vii). Candela : The candela is the luminous intensity, in a given direction, of a source that emits monochromatic radiation of frequency 540 × 1012 hertz and that has a radiant intensity in that direction of 1/683 watt per steradian.

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Uncertainty in Measurement

All scientific measurements involve certain degree of error or uncertainty. Scientific notations, significant figures and dimensional analysis help us in many ways in presenting of data and theoretical calculations.

Scientific Notation

It is an exponential notation in which any number can be represented in the form N x 10n where n is an exponent having positive or negative values and N can vary between 1 to 10. Thus, 232.508 can be written as 2.32508 x 102 in scientific notation.

Precision and Accuracy

Precision refers to the closeness of various measurements for the same quantity. However, accuracy is the agreement of a particular value to the true value of the result. Let the true value of a quantity is 3.9 and its measurements taken by two boys are 3.6 and 3.8. Here 3.8 is more accurate as it is closer to the true value. Similarly 3.85 is more precise than 3.9.

Significant Figures

The total number of digits in measuring of any physical quantity with certainty is called significant figures. There are certain rules for determining the number of significant figures.

  1. All digits are significant except zero in the beginning of a number. For example, in 285 cm, there are three significant figures.

  2. Zeros to the left of the first non-zero digit are not significant if such zeros follow the decimal point. For example, 0.03 has one significant figure.

  3. Zeros to the right of the decimal point are significant. For example, 0.200 g has three significant figures.

  4. Zeros between two non-zero digits are significant. Thus, 2.005 has four significant figures.

  5. Counting the numbers of object, 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 zeroes after placing a decimal i.e., 2 = 2.000000 or 20 = 20.000000.

Notes : In additions or subtractions, the final result should be reported to the same number of decimal places as that of the term with the least number of decimal places.

Laws of Chemical Combination

All chemical reactions take place according to certain laws. These laws are known as laws of chemical combination.

(i). Law of conservation of mass

This law was put forth by Antoine Lavoisier in 1789. According to this, "It states that the total mass of reactants is equal to the total mass of the products".

(ii). Law of constant composition

This law was given by, a French chemist, Joseph Proust. This law states that a chemical compound is always found to be made of same elements combined together in fixed proportion by weight. CO2 can be prepared by number of methods but always 12 g carbon react with 32 g of oxygen.

(iii). Law of multiple proportions

This law was proposed by Dalton in 1803. According to this law, When two elements combine to form two or more chemical compounds, then weight of one of the element which combines with a fixed weight of the other, bears a simple whole number ratio to one another. This is called the law of multiple proportions. For example, The ratio of masses of oxygen in CO and CO2 for fixed mass of carbon (12) is 16 : 32 = 1 : 2

(iv). Law of reciprocal proportions

It states that the ratio of weights of two elements A and B, which combine separately with the fixed weight of a third element C is either same or some simple whole number of the ratio of weights in which A and B combine directly with each other. For example, ratio of masses of carbon and sulphur which combine with the fixed mass (32 parts) of oxygen is 12 : 32 or 3 : 8.

(v). 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 the same temperature and pressure. For example, One volume of hydrogen and one volume of chlorine always combine to form two volumes of HCl gas.

(vi). Avogadro’s Law

In 1811, Avogadro proposed that equal volumes of all gases at the same temperature and pressure should contain equal number of molecules.

Dalton's Atomic Theory

John Dalton in 1808 published “A New System of Chemical Philosophy” in which he proposed atomic theory of matter. The main points of Dalton’s atomic theory are as follows :

  1. Matter is made up of extremely small, indivisible particles called atoms.

  2. Atoms of a given element are identical in all respect, i.e., they possess same size, shape, mass, chemical properties etc.

  3. Atoms of different elements are different in all respects, i.e., they possess different sizes, shapes, masses, chemical properties etc.

  4. Atoms of different elements may combine with each other in a fixed, simple, whole number ratio to form compounds.

  5. Atoms can neither be created nor destroyed in a chemical reaction. Dalton’s theory could explain the laws of chemical combination.

Atomic Mass and Molecular Mass

(i). Atomic Mass

Atomic mass can be defined as a mass of a single atom which is measured in atomic mass unit (amu) or unified mass (u) where,

1 a.m.u. = `1/12`th of the mass of one C-12 atom

(i). 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. Molecular mass expressed in grams is known as gram molecular mass.

Molecular mass of methane,

(CH4) = (12.011 u) + 4 (1.008 u) = 16.043 u

The Mole

One mole is the amount of substance that contains as many as entities as number of atoms in exactly 12.00 g of C-12.

Number of carbon atoms in 12 g of C-12 = 6.022 × 1023

Chemical Formulae

Symbolic representation of compound is called chemical formula. It is of following types :

(i). Empirical Formula

An empirical formula represents the simplest whole number ratio of various atoms present in a compound.

(ii). Molecular Formula

The molecular formula shows the exact number of different types of atoms present in a molecule of a compound.

Relationship between Empirical and Molecular Formula

Molecular formula = (Empirical formula) x n

Measurement of Concentration

The concentration of a solution reflects the relative proportion of solute and solvent present in the solution. The various concentration terms are

1. weight percent (% w/W) = (Weight of solute / Weight of solution) x 100

2. Volume percent (% V/V) = (Volume of solute / Volume of solution) x 100

3. Molality (m) – It is defined as number of moles of solute present in 1 kg of solvent.

m = {Number of moles of solute/Mass of solvent (in kg)} x 100

4. Molarity (M) – It is defined as number of moles of solute present in 1 L of solution.

M = {Number of moles of solute/Volume of solution (in litre)} x 100

5. Mole fraction : Suppose, n is the moles of solute and N is the moles of solvent, then,

(i). Mole fraction of solute (Xsolute) =`\frac{n}{n+N]`

(ii). Mole fraction of solvent (Xsolvent) =`\frac{N}{n+N}`

Xsolute + Xsolvent = 1

6. Normality : It is defined as gram equivalent of solute dissolved in one litre solution.

N = {Gram equivalent of solute / Volume of solution (litre)} x 100

Limiting Reagent

Limiting reagent is the reactant which is completely consumed in a reaction. To estimate the amount of product, limiting reagent should be known.

N2(1 mole) + 3H2(3 mole) ⟶ 2NH3(2 mole)

It means 1 mole of N2 react with 3 mole of H2 to produce 2 mole of NH3.


  1. Matter : Anything that occupies space and has mass.

  2. Element : A pure substance which can neither be decomposed into nor built from simpler substances by any physical or chemical method. It contains only one kind of atoms.

  3. Compound : A pure substance which can be decomposed into simpler substances by some suitable chemical method. It contains only one kind of molecules.

  4. Mixture : A substance obtained by simple mixing of two or more pure substances.

  5. Law of Conservation of Mass : During any physical or chemical change total mass of the products formed is equal to the total mass of the reactants consumed.

  6. Law of Constant Composition : A chemical compound always contains same elements combined together in same proportion of mass.

  7. Law of Multiple Proportions : When two elements combine with each other to form two or more than two compounds then the masses of one of the elements that combine with the fixed mass of the other, bear a simple whole number ratio to one another.

  8. Gay Lussac's law : When gases react with each other they do so in volumes which bear a simple whole number ratio to one another and to the volumes of products, if there are also gases, provided all volumes are measured under similar conditions of temperature and pressure.

  9. Avogadro's Law : Equal volume of all gases under similar conditions contain equal number of molecules.

  10. Atom : The smallest particle of an element that takes part in chemical reactions.

  11. Molecule : The smallest particle of a substance that has independent existence.

  12. Atomicity : The number of atoms in a molecule of the elementary substance.

  13. Unified Mass (u) : One-twelfth of the actual mass of an atom of carbon (C-12). It is equal to 1.66 × 10–27 kg.

  14. Atomic Mass : The average relative mass of an atom of the element as compared with mass of a carbon atom (C-12) taken as 12 u.

  15. Molecular Mass : The average relative mass of a molecule of the substance as compared with mass of an atom of carbon (C-12) taken as 12 u.

  16. Gram Atomic Mass : The mass of 1 mole of atoms (6 × 1023) in g is called gram atomic mass.

  17. Gram Molecular Mass : The mass of 1 gram molecule of compound expressed in grams.

  18. Avogadro's Number (NA) : 6.022 × 1023.

  19. Mole : 6.022 × 1023 specified particles.

  20. Molar Mass : Mass of one mole particles of the substance.

  21. Gram Molecular Volume (G.M.V.) : Volume occupied by one mole molecules of the gaseous substance. Its value is equal to 22.4 L and S.T.P.

  22. Empirical Formula : The formula which gives the simplest whole number ratio of atoms of different elements present in the molecule of the compound. Molecular formula is whole number multiple of empirical formula.

  23. Molarity (M) : Number of moles of solute per litre of solution. Expressed as moles per litre or moles per dm3 or Molar (M).

  24. Molarity changes with change in temperature because volume of the solution changes with change in temperature.

  25. K = °C + 273.15

  26. `°F=9/5(°C)+32`

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