Ncert class 11 chapter 1

Summary of Chapter 1 Chemistry of class 11

1. Basic Concepts of Chemistry

2. Chemistry is called the science of atoms and molecules . Chemistry is the branch of science that studies the composition, properties and interaction in knowing how chemical transformation occur. It plays an important role in daily human needs for food, health care products and other materials aimed at improving the quality of life . Manufacture of life saving drugs and safer alternatives to environmentally hazardous refrigerants are its fair degree of success. What is chemistry?

3. Anything which has mass and occupies space is called matter. Matter Solid ⇌ Liquid ⇌ Gas heat heat cool cool

4. Carbon Dioxide Molecule (CO ₂ ) Water Molecule (H ₂O )

5. Properties of matter and their measurement Physical Properties: properties which can be measured or observed without changing the identity or composition of the substances. E g. : color, odour, etc . Chemical Properties: properties which can be measured or observed when a chemical change occur. Eg. : acidity, basicity, etc.

6. The International system of Units (SI) was established in the 11 th General Conference on Weights and Measurement. NOTE : conversion of units from one system to another is termed dimensional analysis or factor label method or unit factor method .

7. Mass and weight: SI unit is kg Mass: the amount of matter present in a substance. Weight : the force exerted by gravity on an object. Volume: SI unit is m³ For liquids L is used. ( 1L = 1000mL, 1000cm³ = 1dm³ ) Volume: the amount of space that a substance or object occupies, or that is enclosed within a container. Density: SI unit is kg m -3 Density: the amount of mass per unit volume. Temperature: SI unit is K Temperature: the degree or intensity of heat present in a substance or object The other units of temperature are °C and °F . The inter-relations are: °F = 9/5( °C ) + 32 K = °C + 273.15 NOTE : Temperature below 0 °C is possible but in Kelvin scale, negative temperature is not possible .

8. Uncertainty in Measurement The problem of doing simple mathematics with very large or very small numbers is solved by the using scientific notation i.e., exponential notation. N * 10ⁿ where n is an exponent having positive or negative value. Precision: the closeness of various measurement for the same quantity . Accuracy: the agreement of a particular value to the true value of the results . Significant values: meaningless digits which are known with certainty  

9. Rules for determining significant numbers-  All non-zero digits are always significant. Eg . 285 has three significant values . Any zeros between two significant digits are significant Zeroes preceding the first non-zero digit are not significant. A final zero or trailing zeros in the decimal portion ONLY are significant Eg.  .5 00  or .632 000  the zeros are significant         . 00 6  or . 000 96 the zeros are NOT significant Zeroes at the end or right of a number are significant provided they’re on the right side of the decimal. Eg. 100 has one significant figure 100. has three significant figures 100.0 has four significant figures Counting numbers of object have infinite significant figures as there are exact numbers   

10. LAWS OF CHEMICAL COMBINATION LAW OF CONSERVATION OF MASS Matter can neither be created nor destroyed. LAW OF DEFINITE PROPORTION/COMPSITION A given compound always contains exactly the same proportion of element by weight. LAW OF MULTIPLE PROPORTION If two elements 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 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. AVAGADRO’S LAW Equal volumes of gases at same temperature and pressure should contain equal number of molecules.

11. 1) Matter consist of indivisible atoms. 2) All atoms of a given element are identical in properties including mass. 3) Compounds are formed when atoms of different element combine in a fixed ratio. 4) A chemical reaction is a rearrangement of atoms. Dalton's Atomic Theory

12. ATOMIC MASS One atomic mass is defined as the mass exactly equal to one-twelfth the mass of one carbon-12 ( ¹²C ) atom. In this system, ¹²C is assigned a mass exactly 12 atomic mass unit ( amu ) and amasses of all other atoms are given relative to this standard. Atomic mass unit ‘amu’ is replaced by unified mass ‘u’. MOLECULAR MASS Molecular mass is the sum of atomic masses of the elements present in a molecule. Molecular mass = (A₁*N ₁) + (A₂*N ₂) +…. A=atomic mass of each element N=number of its atoms FORMULA MASS Formula mass is the sum of the average atomic masses of all the atoms in the compound’s formula.

13. MOLE CONCEPT AND MOLAR MASSES One mole is the amount of a substance that contains as many particles or entities as there are atoms in exactly 12g of the ¹²C isotope. Knowing one mole of carbon weighs 12g,the number of atoms in it is equal to: = 12g/mol ¹²C . 1.992648 * 10¯²³ g/ ¹²C atom = 6.022 * 10²³ atoms/mol This number is termed Avogadro's constant and denoted by N A. The mass of one mole of a substance in grams is called its molar mass. Molar mass= atomic mass/ molecular mass/ formula mass in u

14. PERCENTAGE COMPOSITION Mass % of an element = mass of that in the compound * 100 molar mass of the compound EMPIRICAL FORMULA FOR MOLECULAR FORMULA An empirical formula is the simplest whole number ratio of various atoms present in a compound. Molecular formula is the exact number of different types of atoms present in the molecule of a compound

15. The word ‘stoichiometry’ is derived from two Greek words — stoicheion (meaning, element) and metron (meaning, measure). Stoichiometry deals with the calculation of masses/volumes of the reactants and the products involved in a chemical reaction. Let us consider the combustion of methane. A balanced equation for this reaction is as given below: CH 4 (g) + 2 O 2 (g) → CO 2 (g) + 2 H 2 O (g) Here , methane and dioxygen are called reactants and carbon dioxide and water are called products . The coefficients 2 for O 2 and H 2 O are called stoichiometric coefficients. Similarly the coefficient for CH 4 and CO 2 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.

16. Thus, according to the above chemical reaction: • One mole of CH 4 (g) reacts with two moles of O 2 (g) to give one mole of CO 2 (g) and two moles of H 2 O(g) • One molecule of CH4(g) reacts with 2 molecules of O 2 (g) to give one molecule of CO2(g) and 2 molecules of H 2 O(g) • 22.7 L of CH 4 (g) reacts with 45.4 L of O 2 (g) to give 22.7 L of CO 2 (g) and 45.4 L of H 2 O(g) • 16 g of CH 4 (g) reacts with 2×32 g of O 2 (g) to give 44 g of CO 2 (g) and 2×18 g of H 2 O (g). CH 4 (g) + 2 O 2 (g) → CO 2 (g) + 2 H 2 O (g) mass  ⇌ moles  ⇌ no . of molecule Mass_ Volume = Density

17. The  limiting reagent  (or  limiting reactant  or limiting agent) in a chemical reaction is a reactant that is totally consumed when the chemical reaction is completed. Many a time, reactions are carried out with the amounts of reactants that are different than the amounts as required by a balanced chemical reaction. In such situations, one reactant is in more amount than the amount required by balanced chemical reaction. The reactant which is present in the least amount gets consumed after sometime and after that further reaction does not take place whatever be the amount of the other reactant. Hence, the reactant, which gets consumed first, limits the amount of product formed and is, therefore, called the limiting reagent.

18. Reactions in Solutions The concentration of a solution or the amount of substance present in its given volume can be expressed in any of the following ways. 1 . Mass per cent or weight per cent (w/w %) Mass per cent= mass of solute * 100 mass of solution 2. Mole fraction If a substance ‘A’ dissolves in substance ‘B’ and their number of moles are n A and n B , respectively, then the mole fractions of A and B are given as: Mole fraction of A= No. of moles of A__ No. of moles of solution = n A . n A + n B Mole fraction of B= No. of moles of B___ No. of moles of solution = n B . n A +n B

19. 3. Molarity It 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. Molarity (M) = No. of moles of solute _ Volume of solution in litres 4. Molality It is defined as the number of moles of solute present in 1 kg of solvent. It is denoted by m. Molality(m) = No. of moles of solute _ Mass of solvent in kg