Chemical Equilibrium PDF

Chemical equilibrium/equilibrium reaction Occurs when the concentrations of the products and reactants in a chemical reaction exhibit no net change over a period of time; The rate of forward reaction equals the rate of backward/reverse reaction Irreversible Reaction Chemical reaction th...

Chemical equilibrium/equilibrium reaction Occurs when the concentrations of the products and reactants in a chemical reaction exhibit no net change over a period of time; The rate of forward reaction equals the rate of backward/reverse reaction Irreversible Reaction Chemical reaction that proceeds in one direction only and goes to completion the reaction continues until one/both reactant(s) is/are used up. Zn + 2HCl ⟶ ZnCl2 + H2 HNO3 + KOH ⟶ KNO3 + H2O  burning of firewood/dry grass/paper. NB: Few chemical reactions proceed in one direction only. Most chemical reactions are reversible at least to some extent. Reversible Reaction Chemical reaction that proceeds in both directions under suitable conditions H2O(l) ⇌ H2O(g) Characteristics of reversible reaction at the start of the reversible process, the reactants proceed towards the formation of products. the reverse process only occurs/takes place when product molecules form. Reversible reactants lead to equilibrium at a time, t.  Chemical equilibrium and the state of dynamic equilibrium condition in which the rate of the forward reaction is exactly balanced by the rate of the backward reaction. occurs when both the forward and backward/reverse reactions are proceeding at the same rate. no net change occurs in the concentration of the reactant/product molecules. Characteristics of dynamic equilibrium energy change in dynamic equilibrium is zero; Concentration of reactants and products is constant;  proceeds in both forward and backward directions/rate of forward reaction equals that of reverse reaction; can either be homogeneous/heterogeneous systems. homogeneous system: all the products and reactants are in the same physical state. 2SO2(g) + O2(g) ⇌ 2SO3(g) heterogeneous system: products and reactants are not in the same physical state. CaCO3(s) ⇌ CaO(s) + CO2(g) Br2(l) ⇌ Br2(g)  Equilibrium constant/equilibrium law (law of mass action) Equilibrium law rate of a chemical reaction is directly proportional to the product of the concentration of reactants at a constant temperature; Consider the general representation of equilibrium reaction: mA + nB ⇌ xC + yD said to be the RDS Rate for forward reaction = Kf[A] m[B] n Rate for reverse reaction = Kr[C] x[D] y At equilibrium, Rate for forward reaction = Rate for reverse reaction Kf[A]m[B]n = Kr[C]x[D]y Kf[A]m[B]n = Kr[C]x[D]y Kf[A]m[B]n Kr[C]x[D]y Kr[A]m[B]n = Kr[A]m[B]n Kf [C]x[D]y = =K=Q Kr [A]m[B]n Equilibrium/reaction quotient, Q contains only concentrations or partial pressures of reacting species; It is the mathematical form of the law of mass action; It is equal to K at constant temperature, meaning system is in equilibrium. Equilibrium constant: constant of proportionality which equals the ratio of the equilibrium concentrations of products to that of reactants which raised to the exponent (power) of its stoichiometric coefficient.  Equilibrium constant for concentration, kC: [C]x[D]y KC = [A]m[B]n = QC Equilibrium constant for partial pressure, kP: From PV = nRT # # P = $RT, $ = [] = c; P = cRT; P ∝ []; (PC)x(PD)y KP = for gaseous species. (PA)m(PB)n PC = partial pressure of C, PD = partial pressure of D, PA = partial pressure of A, PB = partial pressure of B  Relationship between KC and KP From P = cRT PC = [C]RT; PD = [D]RT; PA = [A]RT; PB = [B]RT (PC)x(PD)y ([C]RT)x(D]RT)y KP = ⟹ KP = (PA)m(PB)n ([A]RT)m([B]RT )n [C]x[D]y KP = [A]m[B]n. (RT)(x + y) – (m + n) ; (x + y) – (m + n) = ∆n KP = KC(RT)∆& For pure solids and liquids, the ratio of the amount of substance to volume of substance is a constant (their concentration is constant); They do not appear in KC and KP expressions;  expressions involving aqueous species do not appear in KP because they do not exert pressure. MgCO3(s) ⇌ MgO(s) + CO2(g) KC = [CO2] and KP = PCO2 HF(aq) + H2O(l) ⇌ + F(()) + H3O, (()) F – [H3O+] KC = , no Kp because the reaction does not involve gases. [HF ] General uses of equilibrium constant, K: helps to predict the direction in which a reaction mixture will proceed to achieve equilibrium; helps to calculate equilibrium concentration/partial pressures of products and reactants.  Le Chatelier’s Principle and equilibrium disturbance If a chemical system in equilibrium is subjected to a change in conditions, the system adjusts itself so as to counteract/annul the change; If the conditions of a system in equilibrium changes, the system undergoes a change in order to restore its equilibrium. Equilibrium disturbance: disruption/interruption in an equilibrium system as a result of imposition of external factor(s)/changes in the condition(s) of the system.  Factors/conditions that disturb equilibrium systems: change in concentration (removal or addition of products/reactants); change in temperature; change in pressure/volume; Changes in concentration If the concentration of one of the reacting species present in equilibrium changes without change in any of the other conditions, the position of the equilibrium will have to decrease the concentration of the added substance. change in concentration of substance effect on equilibrium position of reaction: A+B⇌ C+D increase in concentration of A or B conc. of C and D increase; equilibrium shifts to the right/product side decrease in conc. of A or B conc. of C and D decrease; equilibrium shifts to the left/reactant side increase in conc. of C or D conc. of A and B increase; equilibrium shifts to the left/reactant side decrease in conc. of C or D conc. of A and B decrease; equilibrium shifts to the right/product side  Changes in Temperature increasing the temperature of a system in equilibrium favours an endothermic (heat absorbing) reaction; decreasing the temperature favours an exothermic (heat-releasing) reaction; changing the temperature changes the equilibrium constant (K) for a reaction; increasing the temperature of an exothermic reaction makes its K smaller; increasing the temperature of an endothermic reaction makes its K larger and vice versa. Explanation Consider the industry synthesis of ammonia: N2(g) + 3H2(g) ⇌ 2NH3(g) ∆H = – 92 k J mol – 1 [NH3]2 KC = [N2][H2]3 enthalpy change of the reaction is the critical factor; the reaction above is exothermic, increasing temperature shifts the equilibrium to the left/reactant side; the concentration of NH3 (product) will decrease and the conc. of the H2 and N2 (reactants) increase; the numerator (value of product) of the mass action expression becomes larger; this results in smaller reaction quotient;  smaller value of Kc; the reverse process occurs when temperature decrease. Change in pressure changes in pressure/volume (gases) change in pressure affects only volume of a gas; Solids/liquids are much less compressible; when pressure increases (decrease in volume), the equilibrium shifts to the side of the equation that has fewer moles of gas: N2(g) + 3H2(g) ⇌ 2NH3(g) forward reaction favoured; decreasing pressure (increase in volume) shifts equilibrium to the side having more moles of gas: N2(g) + 3H2(g) ⇌ 2NH3(g) backward reaction favoured; when amount of substance of reactants equals amount of substance of products, a change in pressure will have no effect on the equilibrium position of the system: H2(g) + CO2(g) ⇌ H2O(g) + CO(g) Catalysts presence of catalysts has no effect on equilibrium position; the catalyst affects both the forward and reverse reactions equally;  it speeds up forward and backward reaction equally/ to the same degree/extent; the catalyst only helps the system to attain equilibrium faster. Inert Gases inert gas added to a system in equilibrium; concentrations of reactants and products will not change provided this gas does not react with the species in equilibrium; the concentration will continue to satisfy the equilibrium law; the reaction quotient, Q will continue to equals KC;  there will not be any change in the position of the equilibrium. Relation between Q and K and equilibrium position: Equilibrium position is the extent of the reaction in the position of the equilibrium. When Q = K, system is in equilibrium, no shift in equilibrium position; When Q > K, equilibrium shifts to the left; When Q < K, equilibrium shifts to the right.  Significance of K (Kc/Kp) and equilibrium position: value of K depends on only equilibrium concentration of reactants and products. value of K is constant at a given temperature. when Kc = 1, amount of products and reactants are the same; no shift in equilibrium position. when K > 1, large amount of product is present/small amount of reactant is present; equilibrium shifts to product side. If K < 1, only small amount of product is present/large amount of reactant is present; equilibrium shifts to reactant side.