PHYSICS-15-10- 11th (PQRS) SOLUTION

https://docs.google.com/document/d/1P6w4d_jPHSg4pi-fusdDESgv-bViTqr4/edit?usp=sharing&ouid=109474854956598892099&rtpof=true&sd=true Some Basic Concepts of Chemistry Matter and its Nature, Dalton's Atomic Theory; Concepts of Atom, Molecule, Element and Compound; Physical Quantities and their Measurements in Chemistry, Precision and Accuracy, Significant Figures, Units, Dimensional Analysis; Laws of Chemical Combination; Atomic and M olecular M asses, M ole Concept, M olar M ass, Percentage Composition, Empirical and Molecular Formulae; Chemical Equations and Stoichiometry. LAWS OF CHEMICAL COMBINATION There are five laws of chemical combination. Law of Conservation of Mass (Lavoisier 1774). It deals with the mass of reactants and products and states that in a chemical change the total mass of the products is equal to the total mass of the reactants. e.g. C + O2 → CO2 12g + 32g = 44g Law of Constant Composition (Proust 1799). A chemical compound always contains same elements combined together in same proportion by mass e.g H2O prepared from any source contains H and O in the ratio of 1 : 8 by mass. Law of Multiple Proportion (Dalton 1804). When two elements combine to form two or more compounds then the masses of one of which combines with a fixed mass of the other element bears a simple whole number ratio to one another. Law of Reciprocal Proportion (Richter 1792). It states that when two elements combine separately with a fixed mass of the third element then the ratio in which they do so is either the same or whole number multiple of the ratio in which they combine with each other. Gay Lussac’s Law of Combining Volumes. It states that at a given temperature and pressure, when the gases combine they do so in volumes which bear a simple ratio to each other and also to the volume of gaseous product e.g. H2 (g) + Cl2(g) → 2HCl(g) . The ratio of their volumes is 1 : 1 : 2 C H A P T E R CHAPTER INCLUDES Laws of chemical combination Mole concept Measurement of concentration Limiting reagent Equivalent Weight n-factor or valence factor Laws of equivalence Empirical and molecular formula (1L ) (1L ) (2L ) MOLE CONCEPT A mole is nothing but a collection of 6.023 × 1023 particles (atoms or molecules or ions) as well as it is equal to atomic weight (in g), molecular weight (in g) and ionic weight (in g) for atoms, molecules and ions respectively. 1 mole is represented in the forms of atoms, molecules and ions as:- For atoms → 1 gm atom For molecules → 1 gm molecule or 1 gm mole For ions → 1 gm ion Moles can be calculated by the following ways : Number of moles of molecules = Weight of substance (in g) Molecular weight Number of moles of atoms = Number of moles of gases = Weight of substance (in g) Atomic weight Volume of gas at NTP (in litres) 22.4 (1 mole of any gas occupies a volume of 22.4 litres at N.T.P., N.T.P. Corresponds to 0ºC and 1 atm pressure) Number of moles of atoms/molecules/ions = Number of atoms / molecules/ ions Avogadro constant (Avogadro constant is equal to 6.023 × 1023). 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 % w/W (weight percent or Mass percent) x % w/W means that x g solute is present in 100 g of solution. % w/V — x % w/V means that x g of solute is present in 100 ml solution. % v/V — x % v/V means that x ml of solute is present in 100 ml solution. Molality (m) – It is defined as number of moles of solute present in 1 kilogram of solvent. Molarity (M) – It is defined as number of moles of solute present in 1 litre of solution. M = Number of moles of solute Volume of solution (in litre) Suppose, w gram solute is dissolved in V (in ml) solution and molecular weight of solute is m. M = w / m V(in ml)/1000 = w × m 1000 V(in ml) w = Moles of solute. m w ×1000 = m Millimoles of solute. So, millimoles of solute = M × V (in ml) Moles of solute = M × V (in litre). LIMITING REAGENT In the given reaction if number of quantities (either in gm/mole/molecules) if are present with exact co-efficients then it is referred as, reactants are present in exact molar proportions required by chemical equation. However if exact proportion is not present then the one which gets totally consumed is known as limiting reagent (Limiting reagent decides the product quantity for given information). e.g., 2H2 + O2 → 2H2O In above e.g. 2 moles of H2 reacts exactly with 1 mole of O2 to give 2 moles of H2O. If given moles of H2 are 4 moles and that of O2 are 0.5, then 0.5 O2 will act as limiting reagent as it is in minimum amount and product formation is given w.r.t. O2 i.e., 1 mole of H2O. EQUIVALENT WEIGHT Equivalent weight of substance is defined as number of parts by weight of given substance which combines or displaces 1 part by weight of hydrogen (11.2 L of H2 at STP), 8 parts by weight of oxygen (5.6 L at STP), 35.5 parts by weight of chlorine (11.2 L at STP). Atomic weight Equivalent weight of element = Valency Molecular weight of acid Equivalent weight of acids = Basicity Equivalent weight of bases = Equivalent weight of salts = n-FACTOR OR VALENCE FACTOR n-factor is very important for both redox and non redox reactions through which we predict the following two informations: It predicts the molar ratio of the species taking part in reactions i.e. reactants. The reciprocal of n-factor's ratio of the reactants is the molar ratio of the reactants. For example : If X (having n-factor = a) reacts with Y (having n-factor = b) then its n-factor's ratio is a : b, so molar ratio of X to Y is b : a. It can be represented as bX + (nf =a) aY (nf =b) ⎯⎯→Pr oducts Equivalent weight = LAWS OF EQUIVALENCE Molecular weight n − factor or Atomic weight n − factor According to law of equivalence, for each and every reactant and product, Equivalents of each reactant reacted = Equivalents of each product formed. For example : Suppose, the reaction is taking place as under A + B → C + D. Then according to law of equivalence, Equivalents of A reacted = Equivalents of B reacted = Equivalents of C produced = Equivalents of D produced Equivalents of any substance = Weight of substance (in g) Equivalent weight = Normality (N) × Volume (V) (In litre) Normality (N) = n-factor × Molarity (M) Normality and molarity are temperature dependent. Since on changing the temperature, the volume of solution changes, so normality and molarity change. EMPIRICAL AND MOLECULAR FORMULA Empirical Formula of a compound is the simplest whole number ratio of the atoms of elements constituting its one molecule. The sum of atomic masses of the atoms representing empirical formula is called empirical formula mass. Molecular Formula of a compound shows the actual number of the atoms of the elements present in its one molecule. The sum of atomic masses of the atoms representing molecule is called molecular mass. Relationship between Empirical Formula and Molecular Formula Molecular formula = n × empirical formula where n is a simple whole number having values of 1, 2, 3... etc. Also, n = Molecular formula mass/Empirical formula mass. ❑ ❑ ❑