Synopsis of STOICHIOMETRY - LAWS OF CHEMICAL COMBINATION • The mass relation between the reactants and products in a chemical reaction is called stoichiometry. • There are four important laws of chemical combinations • THE LAW OF CONSERVATION OF MASS: • This law was proposed by Lavoisier in 1789 by carrying several experiments. • The law states that matter can neither be created nor destroyed during a chemical change. • The law may also be stated as the total mass of the products formed during a chemical change is exactly equal to the total mass of the reactants. • Weighed amounts of solid and solid KI are dissolved in water separately and their solutions are mixed. The following reaction takes place • Total mass of is equal to the total mass of . • LAW OF DEFINITE PROPORTIONS: • Proposed by Proust. Verified by Stress and Richards. • It is also known as Law of constant proportions. • A given compound always contains the same elements combined in a fixed proportions by weight. • What ever the method a compound is prepared, it contains the same elements combined in a fixed ratio by weight • Eg: CO2 can be prepared by many ways i.e., by combining of carbon with oxygen or by heating lime stone etc., but what ever the method CO2 is prepared; The ratio of carbon and oxygen by mass is 12 : 32 = 3 : 8 • LAW OF MULTIPLE PROPORTIONS: • Proposed by Dalton. Verified by Berzelius. • If two elements chemically combine to give two or more compounds, then the weight of one element which combines with the fixed weight of the other element in those compound bear a simple multiple ratio to one another. • Eg: Nitrogen forms the oxides; N2 O, NO, N2O3, NO2, N2O5 • In these compounds 28 gm of Nitrogen combines with 16, 32, 48, 64, 80 gm of oxygen respectively. The weight of oxygen in these compounds are in the ratio 16:32:48:64:80 or 1:2 : 3 : 4 : 5 a simple multiple ratio. LAW OF RECIPROCAL PROPORTIONS: • This law was proposed by Richter (1792) which states as “when two elements combine separately with a fixed mass of a third element, then the ratio of their masses in which they do so is either same or some whole number multiple of the ratio in which they combine with each other. • GAY-LUSSAC’S LAW OF COMBINING VOLUMES • According to this law gases combine in the simple whole number ratio of their volumes under similar conditions of temperature and pressure. If products are also gases, the simple whole number ratio also extends to the products. • Eg: • Under similar conditions, 2 lts of Hydrogen combines with 1lt of oxygen to give 2 lts of water vapour. • It is applicable only to gaseous reaction. • Law of combining volumes can be derived from Law of defininte proportions when expressed in terms of volumes. AVOGADRO’S LAW: • At the same T, P equal volumes of all gases contain equal number of moles or molecules. • No. of molecules = no. of moles × N • The study of mass relation or quantity relation between reactants and products is called Stoichiometry. • According to law of conservation of mass, proposed by Lavoisier, the mass of reactants should be equal to the mass of products. • The balanced chemical equation which gives correct relation between reactants and products is called stoichiometric equation. • The definite quantities which are involved in chemi-cal reaction and present in balanced stoichiometric equation are called stoichiometric quantities. 1) H2 + O2 → H2O 2g + 32g → 18g (x) 2) 2H2 + O2 → 2 H2O 4g + 32g → 36 g Equation 1) is not stoichiometric equation and will not obey law of conservation of mass, because masses are not stoichiometric quantities Equation 2) is stoichiometric equation and will obey law of conservation of mass, because masses are stoichiometric quantities. Isotopic abundance:- It is the percentage availability of an Isotope in the nature. • The atomic wt of Cl is fractional i.e., 35.5 because it has two isotopes and both have significant abundance 35Cl is available by 75% and 37Cl is available by 25% ∴ Atomic weight of Cl (average) = 35.5 100 75 35 25 37 = × + × Though many elements have isotopes, their atomic weights are whole numbers, because of negligible abundance of their isotopes. Protium deuterium tritium Atomic weight = 1 Atomic weights:- The relative atomic weights of elements are expressed in a.m.u. 1 amu = 1.66 × 10–24 g. weight of one H-atom = 1 amu weight of 1/16 of O-atom = 1 amu weight of 1/12 of C-atom = 1 amu The atomic weights are measured by C-12 scale. The atomic weight of an element is the number of times heavier when compared to 1/12th of C atom. Atomic weight = 1/12 of carbon atomic weight weight of atom of element Since, the atomic weight of an element is ratio, it has no units. Molecular weights:- Even the molecular weights can be measured with the help of C-12 scale. The molecular weight of a substance is defined as the number of times when compared to 1/12th of carbon atom. molecular weight = weight of 1/12th of carbon weight of one molecule Thus, relative atomic weight and molecular weights are expressed in amu, but they dont have units. Eg.: 1. Atomic weight of Hydrogen = 1 amu. Actual weight of Hydrogen = 1.66 × 10–24 g Actual weight of 10 H-atoms= 10×1.66×10– 24 g 2. Atomic weight of oxygen = 16 amu Actual weight of oxygen = 16×1.66×10–24g Actual weight of 100 oxygen atoms= 100×16×1.66×10–24g Gram atom (Gram atomic weight): If atomic weight is expressed in grams it is called gram atomic weight or gram atom. 1 gram atom of Hydrogen = 1 g 2 gram atom of Hydrogen = 2 g 1 gram atom of Oxygen = 16 g 2 gram atom of Oxygen = 32 g 4 gram atom of Oxygen = 64 g Hydrogen : Atomic weight = 1 Gram Atomic weight = 1 g Oxygen : Atomic weight = 16 Gram Atomic weight = 16 g Gram molecule (Gram molecular weight or Gram mole): If molecular weight is expressed in grams, it is called gram molecule or gram molecular weight. Ex.: Hydrogen : Molecular weight = 2 Gram Molecular weight = 2 g Oxygen : Molecular weight = 32 Gram Molecular weight = 32 g CO2 : Molecular weight = 44 Gram Molecular weight = 44g 64g of Oxygen = 2 gram molecule of O2 22g of CO2 = 1/2 gram molecule of CO2 360g of H2O = 20 gram molecule of H2O Mole Concept:- The amount of substance which contains Avagadros number of particles is called mole. (or) Mole is the amount of substance containing as many particles as the number of atoms in 12g of carbon. 1 mole of Hydrogen = 6.023 × 1023 molecules 1 gm molecule weight = 2 g. 1 mole of Carbon = 6.023 × 1023 atoms 1 gm atom weight = 12 g 1 mole of Carbondioxide = 6.023 ×1023 molecules 1 gm molecule weight = 44 g. 1 mole of Sodium = 6.023 × 1023 atoms 1 gm atom weight = 23 g. 1 mole of Sulphur dioxide = 6.023×1023 molecules 1 gm molecule weight = 64 g. 1 mole of Oxygen = 6.023 × 1023 molecules 1 gm molecule weight = 32 g. 1 mole of H2SO4 = 6.023 × 1023 molecules 1 gm molecule weight = 98 g. Gram Molar volume (or) Gram Molecular Volume (GMV): It is the volume occupied by one mole of gas at STP i.e. 22.4 lit. 44 g of CO2 occupies 22.4 lit at STP 11 g of CO2 occupies 5.6 lit at STP 8 g of CH4 occupies 11.2 lit at STP 4 g of He occupies 22.4 lit at STP 142 g of Cl2 occupies 2 × 22.4 lit at STP Weight of 5.6 lit of Ethane at STP is 7.5 g Weight of 44.8 lit of SO2 at STP is 128 g. 22.4 lit of any gas at STP contains 6×1023 molecules 11.2 lit of Chlorine contains 3 × 1023 molecules 5.6 lit of Hydrogen contains 1.5 × 1023 molecules 3 × 1023 molecules of CO2 will occupy 11.2 lit 6 × 1023 molecules of O2 will occupy 22.4 lit 6 × 1024 atoms of O2 will occupy 224 lit 3 × 1024 atoms of O2 occupies a volume of 112 lit.
Posted on: Fri, 29 Nov 2013 06:48:36 +0000
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