Chemical Equilibrium

Reversible Reactions Reversible reactions are those in which the products can react to re-form reactants. Ex: 2HgO (s) + heat > 2Hg (l) + O 2 (g) 2Hg (l) + O 2 (g) > 2HgO (s) + heat Chemical Equilibrium: When the rate of the forward reaction = the rate of the reverse reaction AND the concentrations of product and reactants remain the same. Freezing Ice? Cooking an egg? Melting water?

Equilibrium If product is favored, the reaction is represented as: HBr (aq) + H 2 O(l) H 3 O + (aq) + Br - (aq) If reactant is favored, the reaction is represented as: H 2 CO 3 (aq) + H 2 O (l) H 3 O + (aq) + HCO 3 - (aq) If concentrations of reactants = products: H2SO3 (aq) + H2O (l) H3O+ (aq) + HSO3- (aq) The extent to which reactants ----> product is determined from the equilibrium constant > < > < > <

The Equilibrium Expression Since the concentrations of reactants and products after equilibrium has been established remain constant, so does the ratio of reactant to product. Ex: nA(aq)+ mB(aq) > xC(aq) + yD(aq) K = [C] x [D] y If K > 1, product is favored If K < 1, reactant is favored If K = 1, reactant = product [A] n [B] m K is determined experimentally K is dependent on temperature No solids or liquid included in the equilibrium expression

Question What is the equilibrium expression for the following reaction at equilibrium? H 2 (g) + I 2 (g) HI (g) K = [HI] 2 The equilibrium expression can be used to calculate concentrations of reactants and products at equilibrium. Let’s try… > < [H 2 ][I 2 ]

Predicting Direction of a Shift in Equilbrium Le Chatelier’s Principle: If a stress is applied to a system at equilbrium, the system will react in a way to minimize that stress. Changes in pressure: –Affect gases only! Changes in temperature Changes in concentration: –Solids and liquids not affected. - K is not affected.

Reactions that Go to Completion In these reactions, ions are almost completely removed from solution. Formation of a gas –Ex: H 2 CO 3 (aq)  H 2 0 (l) + CO 2 (g) Formation of a precipitate –Ex: NaCl(aq) + AgNO 3 (aq)  NaNO 3 + AgCl (s) Formation of a slightly ionized product (acids and bases) Ex: HCl(aq) + NaOH(aq)  NaCl(s) + H 2 O(l)

Common Ion Effect When the addition of an ion common to 2 solutes brings about precipitation. Ex: NaCl (s) Na + (aq) + Cl - (aq) If I now add HCl to this solution, which ion will be affected? HCl(g) + H 2 O(l) -----> H 3 O + (aq) + Cl - (aq) Equilibrium will shift to the left (production of Solid NaCl) > <

Equilibria of Acids, Bases and Salts Ionization of a Weak Acid CH 3 COOH + H 2 O H 3 O + + CH 3 COO - K = [H 3 O + ][CH 3 COO - ]/[CH 3 COOH][H 2 O] We can assume that the [H 2 O] remains constant. K[H 2 O] = [H 3 O + ][CH 3 COO - ] = K a > < [CH 3 COOH] Acid Ionization Constant

KaKa Constant at a given temperature. For weak acids, K a is < 1. To determine the K a, the [ ] of involved reactants/products must be known. [ ] are found by measuring pH. K b used for bases.

Buffers Resist changes in pH Weak acid + its salt (or weak base + salt) Salts of a weak acid can combine with the extra H 3 O + that is produced. –Ex: CH 3 COOH/NaCH 3 COO solution –CH 3 COO - + H 3 O > CH 3 COOH + H 2 O Buffered SolutionAcidNo pH change ―CH 3 COO - + H + + OH -  H 2 0 Buffered SolutionBase first ―CH 3 COOH + H 2 0  CH 3 COO - + H 3 O + second Offsets original base

Hydrolysis of Salts Salts formed during neutralization Hydrolysis occurs when water molecules combine with ions of a dissolved salt. Can form acidic, basic or neutral solutions: –Strong Acid + Strong Base -----> Neutral –Strong Acid + Weak Base -----> Acidic –Strong Base + Weak Acid -----> Basic –Weak Base + Weak Acid > Unpredictable

Solubility Equilibrium Used to predict whether precipitation occurs when solutions are combined. Soluble (>1 g/100 g water) Ex: AgCl (s) -----> Ag + (aq) + Cl - (aq) K sp = [Ag + ][Cl - ] = Solubility Product Constant –If K sp > ion product, the substance DISSOLVES ( unsaturated- no solid present) –If K sp < ion product, there is PRECIPITATION − Pg. 579 Table 18-3