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

https://docs.google.com/document/d/1kGcRAQxf6nss2IYmI6Fqxl2GzretqmNC/edit?usp=sharing&ouid=109474854956598892099&rtpof=true&sd=true physical processes : Solid -liquid, liquid - gas and solid - gas equilibria, Henry’s law, general characteristics of equilibrium involving physical processes. Equilibria involving chemical processes : Law of chemical equilibrium, equilibrium constants (K p and Kc ) and their significance, significance of ΔG and ΔGº in chemical equilibria, factors affecting equilibrium concentration, pressure, temperature, effect of catalyst; Le Chatelier’s principle. Ionic equilibrium : Weak and strong electrolytes, ionization of electrolytes, various concepts of acids and bases (Arrhenius, Bronsted - Lowry and Lewis) and their ionization, acid - base equilibria (including multistage ionization) and ionization constants, ionization of water, pH scale, common ion effect, hydrolysis of salts and pH of their solutions, solubility of sparingly soluble salts and solubility products, buffer solutions. CHEMICAL EQUILIBRIA It is established in chemical reaction Ex. N2 + 3H2 2NH3(g) N2(g) + O2(g) 2NO(g) The state of equilibrium is a state in which the measurable properties of the system do not undergo any noticeable change under a particular set of conditions. LAW OF MASS ACTION According to Guldberg and Waage, “The rate of a chemical reaction is directly proportional to the product of active masses of the reactants”. Active mass of reactant = [mole L–1] Since, active mass = Concentration of reactant × Activity coefficient. = Molarity × γ For dilute solution; γ = 1 EQUILIBRIUM CONSTANT (K) IN TERMS OF CONCENTRATION AND PRESSURE (KC & KP) CHAPTER INCLUDES Chemical Equilibria Law of Mass Action Equilibrium Constant (K) Relationship between K p and Kc Le–Chatelier's Principle Ionic Equilibrium Acids–Bases pH scale Buffer Solutions Salt Hydrolysis KC = Molar concentration of product or multiplication of molar concentrations of products with stoichiometric quantity as a power Molar concentration of reactant or multiplication of molar concentrations of reactants with stoichiometric quantity as a power Solubility Product KP = Partial pressure of gaseous product or multiplication of partial pressures of gaseous products with stoichiometric quantity as a power Partial pressure of gaseous reactant or multiplication of partial pressures Note : For liquids and solids, the partial pressure is taken to be unity. RELATIONSHIP BETWEEN KP AND KC K = K (RT)Δng where Δ n g = np – nr = no of moles of gaseous product – no. of moles of gaseous reactant. if Δng = 0; Kp = Kc ; Example : N2(g) + O2(g) 2NO(g). if Δng > 0; Kp > Kc ; Example : PCl5(g) PCl3(g) + Cl2(g) if Δng < 0; Kp < Kc ; Unit of Kp and Kc Unit of Kp = (atm)Δng Example : N2 (g) + 3H2 (g) 2NH3(g) Unit of Kc = (mol L−1)Δng If Δng = 0, no unit of Kp and Kc Relation between degree of dissociation ( α ) and Vapour density For equilibrium reaction PCl5 (g) PCl3 (g) + Cl2 (g) N2O4(g) 2NO2(g) For (i) & (ii) n = 2 D = vapour density of the gas before dissociation = molecular weight/2 d = vapour density of equilibrium mixture n = total no. of moles obtained after dissociation from 1 mole of dissociating molecule. Relation between Equilibrium Constant (K) and Standard free energy Change (ΔG°) LE-CHATELIER’S PRINCIPLE It states that, if a system in equilibrium is disturbed by any external agency such as pressure, temperature, concentration etc. then equilibrium will get shifted to counter balance the effect of that disturbance. Factors affecting Equilibrium Concentration Addition of any reactant or removal of products leads to forward reaction or vice-versa. Temperature In an endothermic reaction (ΔH= + ve) increase in temperature favours the forward reaction, while decrease in temperature favours backward reaction. For exothermic reaction (ΔH= − ve) increase in temperature favours backward reaction, while decrease in temperature favours forward reaction. Pressure Effect of pressure is mainly applicable to gaseous reactions, since liquids & solids are incompressible in nature. If Δng = 0 Δng > 0 pressure has no effect on equilibrium constant. then on increasing pressure, equilibrium will get shifted in the backward direction. Δng < 0 then on increasing pressure, equilibrium will shift towards forward direction. Catalyst A catalyst has no effect on state of equilibrium but it enables the state of equilibrium to reach very quickly. Inert gas The introduction of inert gas to any equilibrium is visualized under the condition of constant volume and constant pressure. at constant volume If Δng = 0 Δng > 0 Δng < 0 no effect on equilibrium on addition of inert gas at constant pressure If Δng = 0 no effect on equilibrium Equilibrium will get shifted in that direction where no. of moles are more. if Δng > 0⎪⎫ In forward direction ⎬ Δng < 0⎪ In backward direction Example : Δng > 0 PCl5 (g) PCl3 (g) + Cl2 (g) Example: Δng < 0 2SO2(g) +O2 (g) 2SO3 (g) IONIC EQUILIBRIUM Dissociation of weak acids or weak bases and Ostwald dilution law When weak acid or weak base is dissolved in aqueous medium equilibrium exists between dissociated ions and undissociated molecules. CH3COOH CH3COO– + H+ Initially C 0 0 At equilibrium C (1-α) Cα Cα where α the degree of dissociation ∴ Ka = Cα ⋅ Cα = Cα2 C(a − α) or [If α << 1] and Similarly for weak base or ACIDS AND BASES Arrhenius Concept : An acid is a substance which furnishes H+ ions in aqueous solution HCI H+ + CI– strong acid CH3COOH CH3COO– + H+ weak acid A base is a substance which furnishes OH– ions in aqueous solution NaOH Na+ + OH– strong base NH4OH NH + + OH– weak base The strength of an acid depends upon its tendency to furnish H+ or OH– ions in solution. Bronsted – Lowry Concept : An acid is a proton donor while base is a proton acceptor. NH3 + H2O NH + (aq) + OH– (aq) base1 acid2 acid1 base2 A pair of acids and bases which differ by a proton is known as conjugate pair of acid and base. Stronger is the acid, weaker is its conjugate base and vice versa. The strength of acids (or bases) also depends upon the tendency of base (or acid) which accept proton. Lewis Concept : An acid is an electron pair acceptor. (c) A base is an electron pair donor. EXPRESSING HYDRONIUM ION CONCENTRATION–pH SCALE pH is A scale for measuring acidity of a solution. Mathematically pH = –log [H+] = –log [H3O+] Similarly pOH is defined as A scale for measuring basicity of a solution pOH = – log [OH–] BUFFER SOLUTIONS A Solution whose pH remains constant either on keeping for a longer period or by the addition of small quantity of some external electrolyte (i.e. an acid or a base) is known as Buffer solution. Types of Buffer Solutions Acidic buffer : A mixture of weak acid and its salt with a strong base form acidic buffer. e.g. CH3COOH + CH3COONa HCOOH + HCOONa C6H5COOH +C6H5COONa Basic buffer : It is a mixture of Ex. (1) NH4OH + NH4Cl Cu(OH)2 + CuCl2 Mg(OH)2 + MgSO4 ∴ pH= 14 – pOH ⎛ = 14 – ⎜pKb + log ⎝ [Salt] ⎞ ⎟ [base] ⎠ ⎛ pH = pKw – ⎜pKb + log [Salt] ⎞ ⎟ [base] ⎠ SALT HYDROLYSIS Consider the salt BA, which on hydrolysis will give acid and base inside the aqueous solution. BA + H2O BOH + HA Salt made up of strong acid, strong base will not hydrolyse it will simply ionise. And pH of aqueous solution will be neutral. The pH of aqueous solution is independent with dilution i.e., pH = 7. e.g., NaCl, Na2SO4. Salt made up strong acid and weak base will hydrolyse. e.g., NH4Cl. pH = 1 2 [pKw – pKb – logC] Salt made up of weak acid and strong base. e.g., CH3COONa. 1 pH = 2 [pKw + pKb + logC] Salt made up of weak acid and weak base. The pH of these salt solution depends on Ka and Kb but independent with concentration. e.g., CH3COONH4. 1 pH = 2 [pKw – pKb + pKa] SOLU When a sparingly soluble salt is dissolved in water they form saturated solution but concentration of salt is very low. Therefore in saturated solution of such electrolytes, solid electrolyte is in equilibrium with the ions as represented below: AgCl KH Note : KH = Ch2, h = C , where h = degree of hydrolysis. BILITY PRODUCT Ag+ + Cl– Applying Law of chemical [Ag+ ] [Cl− ] K = [AgCl] K[AgCl] = [Ag+] [Cl–] K sp = [Ag+ ] [Cl− ] where Ksp is a constant known as solubility product. It remains constant at constant temperature for a given salt and defined as the product of ionic concentration of a sparingly soluble electrolyte in a saturated solution. if Ksp = I.P. (Ionic product) then solution is said to be saturated if Ksp > I.P. then solution is said to be unsaturated if Ksp < I.P. then solution is said to be supersaturated (Condition for precipitation) Relation between Solubility and Solubility Product For binary electrolyte; Ksp = s2 (AgCl, AgBr) for ternary electrolyte; Ksp = 4s3 (CaF2, BaCl2) for Quarternary Electrolyte; Ksp = 27s4 ❑ ❑ ❑