Mole Concept Notes

Page # 1 MOLE CONCEPT MOLE CONCEPT 1. SIGNIFICANT FIGURES (A) Every scientific observation involves some degree of uncertainity depending upon the limitation of instrument. To repr esent scientific data, role of significant figures has its own importance. (B) Significant figures are equal to the number of digits in numbers with last digit uncertain and rest all are certain digits i.e. all the digits of datum including the uncertain one , are called significant figures. (C) Rules for determination significant figure: (i) All non zero digits are significant. Example : 3.14 has three significant figures (ii) The zeros to the right of the decimal point are significant. Example : 3.0 has two significant figures. (iii) The zeros to the left of the first non zero digit in a number are not significant. Example : 0.02 has one significant figure. (iv) The zeros between two non zero digits are also significant. Example : 6.01 has three significant figures. (v) Exponential form : N × 10n. Where N show the significant figure. Example : 1.86 × 104 has three significant figure. (vi) Rounding off the uncertain digit : (a) If the left most digit to be rounded off is more than 5 , the preceding number is increased by one. Example : 2.16 is rounded to 2.2 (b) If the left most digit to be rounded off is less than 5 , the preceding number is retained. Example : 2.14 is rounded off to 2.1 (c) If the left most digit to be rounded off is equal to 5 , the preceding number is not changed if it is even and increased by one if it is odd. Example : 3.25 is rounded off to 3.2 2.35 is round off to 2.4 2. TYPES OF MIXTURE 2.1Heterogenous mixture A mixture in which the different constituents are not distributed uniformly is known as heterogenous mixture. eg Water 2.2Homogenous mixture A mixture in which the different constituents are uniformly distributed is known as homogenous mixture. eg. O2, N2 etc. 3. LAWS OF CHEMICAL COMBINATION 3.1Law of conservation of mass-[Lavoisier, 1744] (a) According to this law , matter is neither created nor destroyed in the course of chemical reaction although it may change from one form to other (b) This law contradicts nuclear reactions where Einestein equation is applicable (c) According to this law , sum of the masses of product formed is always equal to the sum of the masses of the reactant undergone change Example : H2 + Cl2  2HCl 2gm 71 gm 2 [ 1 + 35.5]  73 gm  73 gm 3.2Law of definite proportion [Proust, 1799] (a) According to the law , the composition of a compound always remains a constant i.e. the ratio of weights of different elements in a compound ; no matter by whatever method , it is prepared or obtained from different sources, remains always a constant Example : In H2O ratio of weight = 1 : 8 In CO2 ratio of weight = 3 : 8 3.3Law of multiple proportion [John Dalton, 1804] According to this law, when two elements A and B combine to form more than one chemical compound then different weights of A , which combine with a fixed weight of B , are in a proportion of simple whole number MATTER Homogenous (Composition uniform) Heterogeneous (Composition nonuniform) Mixture Elements (only one type of atoms) Compounds * only one type of molecules * Different type of atoms but composition is uniform * Different type molecules of composition of them is not fixed and can be sometimes separated Page # 2 MOLE CONCEPT Example : CO & CO2 12 : 16 & 12 : 32 ratio = 16 : 32 = 1 : 2 3.4Law of reciprocal proportions [Ritche,1792-94] When two elements combines separately with third element and form different types of molecules, their combining ratio is directly reciprocated if they combine directly Example : or Vapour density = Weight of one molecule of Gas Weight of one atom of hydrogen×2 or Vapour density = 2 (5) It helps to determine molar volume Molecular weight of the gas = 2 × vapour density = 2 × = 2 × C with H form methane and with O form CO2. In CH4 , 12 grams of C reacts with 4 grams of H whereas in CO2 12 gram of C reacts with 32 grams of O. Therefore when H combines with O they should combine in the ratio of 4 : 32 (i.e. = 1 : 8) or in simple multiple of it. The same is found to be true in H2O molecule. The ratio of weights of H and O in Water is 1 : 8 3.5 Gay-Lussac’s [1808] law of combining volumes This law states that under similar conditions of pressure and temperature, volume ratio of gases is always in terms of simple integers. Ex. N2(g) + 3H2(g)  2NH3(g) vol. ratio 1 : 3 : 2 4. AVOGADRO'S HYPOTHESIS According to this under similar conditions of pressure and temperature , equal volumes of gases contain equal number of molecules. 4.1Salient features of Avogadro's hypothesis (1) It has removed the anomaly between Dalton's atomic theory and Gay Lussac's law of volume by making a clear distinction in between atoms and molecules (2) It reveals that common elementary gases like hydrogen , nitrogen , oxygen etc. are diatomic (3) It provides a method to determine the atomic weights of gaseous elements (4) It provides a relationship between vapour density and molecular weight of substances Vapour density Volume of definite amount of Gas = Volume of same amount of Hydrogen or Vapour density = Weight ofn molecules of Gas Weight of n molecules of Hydrogen Molecular weight Weight of 1litre of the Gas at S.T.P Weight of 1litre of Hydrogen at S.T.P Weight of 1 litre of the Gas at S.T.P 0.089 gm 2 = × Weight of 1 litre of the gas at S.T.P.. 0.089 = 22.4 × Weight of 1 litre of gas at S.T.P. = Weight of 22.4 litre of the gas at S.T.P 5. ATOM, MOLECULES AND MOLECULAR FORMULA Atom: It is the smallest particle of an element that takes part in a chemical reaction and not capable of independent existence. Molecule : It is the smallest particle of matter which is capable of independent existence. A molecule is generaly an assembly of two or more tightly bonded atoms. Homoatomic molecules : Molecules of an element contain one type of atoms. eg. O2, Cl2 etc. Heteroatomic molecules : Molecules of compounds contain more that one type of atom. eg. H2O, HCl etc 5.1Atomic mass scale (A) Oxygen as standard : The standard reference for atomic weight may be oxygen with an assigned value of 16. Atomic weight of an element = Weight of 1 atom of element 1/16 × Weight of 1 atom of oxygen (B) Carbon as standard : The modern reference standard for atomic weight is carbon isotope of mass number 12. Atomic weight of an element = Weight of 1 atom of the element 1/12 × Weight of 1 atom of C -12 IMPORTANT POINTS (1) Atomic weight is not a weight but a number. (2) Atomic weight is not absolute but relative to the weight of the standard reference element C-12