Acid-Base Equilibria

Unique nature of water and its solutions Acid/Base Equilibria (GC, BC) • Bronsted-Lowry definition of acid, base o BL acid: H+ donor o BL base: H+ acceptor • Ionization of water o Kw, its approximate value (Kw = [H+][OH-] = 10-14 at 25oC, 1 atm) ▪ Kw is constant at a given temperature ▪ Addition of acid does not change Kw • It's dependent on both H and OH concentrations o Definition of pH: pH of pure water • Conjugate acids and bases (e.g. NH4+ and NH3) o Stability of conjugate acid/base dictates strength of original acid/base • Strong acids and bases (e.g. nitric, sulfuric) • Weak acids and bases (e.g. acetic, benzoic) o Dissociation of weak acids and bases with or without added salt ▪ Dissociation of weak acids in water cause 3 reactions to occur: • HA + H2O → A- + H3O+ • A + H2O → HA + OH• H2O + H2O → H3O+ + OH▪ Many equilibrium shifts occur as a result of adding a weak acid • 1) Addition of weak acids shifts equation of 1st reaction to the right, increasing H3O+ • 2) Addition of weak acid shifts equation of 2nd reaction to the left, decreasing OHo Since the amount of CB for a weak acid is usually low, OH isn't reduced by much ▪ Net results of adding a weak acid • [H3O+] increases significantly, [OH- decreases insignificantly] o Hydrolysis of salts of weak acids or bases o Calculation of pH of solutions of salts of weak acids or bases ▪ 1) Find Ka of the acid (usually given to you) ▪ 2) Solve using the eq. constant. Set it equal to concentration of the products of dissociation (A- and H3O+, which concentrations "x" as we don't know how much weak acid dissociated) over concentration of reactants (HA, with a concentration of original concentration minus x) ▪ 3) Remove x from denominator (it's probably hella small) and solve • Equilibrium constants Ka and Kb: pKa, pKb o Ka ▪ Dissociation constant of an acid in water ▪ HA + H2O → H3O+ + A- o Kb ▪ ▪ ▪ Dissociation constant of base in water B + H2O → HB+ + OH- ▪ o KaKb = Kw o Difference between Ka/b and Ksp ▪ ▪ • Ka/b measure the tendency of an acid/base to dissociate (lose or gain a proton) in solution Ksp measures a solid's solubility (it's tendency to go to the aqueous phase) • Ksp implies a phase change while Ka/b don't necessarily Buffers o Salts ▪ Ionic compounds that dissolve in water ▪ Conjugates • Species that form after the removal of a H+ or the addition of OH (if H+ can't be removed from a salt) o Definition and concepts (common buffer systems) ▪ Made from equal and copious amounts of a weak acid and it's CB • pKa of that weak acid should be close to the pH that you wish to buffer the solution at ▪ Addition of acid would be offset by the copious amounts of base already present ▪ pH of buffered solution can be found by Henderson-Hasselbach equation • pH = pKa (of weak acid) + log[A-]/[HA] o Influence on titration curves Ions in Solutions (GC, BC) • Solution formation o 3 steps ▪ 1) Breaking of intermolecular bonds b/w solute molecules (endothermic) ▪ 2) Breaking of intermolecular bonds b/w/ solvent molecules (endothermic) ▪ 3) Forming of intermolecular bonds b/w solvent and solute (exothermic) o Thermodynamics ▪ For MCAT solutions, enthalpy change = internal energy change ▪ Solution formation usually leads to increased entropy since combined mixture is more disordered than separate pure substances • Exception: when a gas dissolves in a liquid or solid. In this case, entropy change is negative o Heat of solution: total heat absorbed or released when something is dissolved in solution ▪ Positive Hsol = formation of weaker intermolecular bonds • Heat absorbed when old intermolecular bonds in solute/solvent were more stable than those created when solvent and solute react ▪ Negative Hsol = formation of stronger intermolecular bonds • Heat released when new intermolecular bonds formed in reaction of solute/solvent are more stable than old ones holding solute/solvent together • Anion, cation: common names, formulas and charges for familiar ions (e.g. NH4+ ammonium, PO43phosphate, SO42- sulfate) • Hydration, the hydronium ion • Amphoteric: a substance that can act as either an acid or base Solubility (GC) • Units of concentration (e.g. molarity) • Solubility product constant; the equilibrium expression of Ksp o Ksp = equilibrium of a solvation reaction ▪ Includes products only • Leaves out solids and pure liquids (usually on reactant side) ▪ Measure of how inclined compound is to dissolve • o In an ideally dilute solution (high solvent to solute ration), solvent obeys Raoult's law and solute obeys Henry's law Common-ion effect, its use in laboratory separations o Common-ion effect ▪ Happens when a soluble compound containing one of the ions in the Ksp is dissolved in the solution ▪ More of the common ion on the right side → equilibrium pushes reaction leftwards → decrease in solubility ▪ Does NOT affect value of Ksp ▪ Does NOT affect value of saturated solution (no leverage to shift equilibrium in, added salt containing common-ion can't dissociate) o Complex ion formation o Complex ions and solubility ▪ Greater size or charge of ionic compounds tends to make them less likely to dissolve o Solubility and pH ▪ Acid removes hydroxide from solution • More acid → less hydroxide concentration → equilibrium shifts towards side with hydroxide Titration (GC) • Two reasons for titration o 1) To find the concentration of a substance o 2) To find the pKa or pKb and the Ka or the Kb • Indicators o Used to find the equivalence point in a titration o Usually a weak acid whose CB is a different color ▪ Color can change from acid → CB or from CB → acid • Therefore, depending on which species you start at, there are two pH's where color change can occur • Range of an indicator: pH values of two points of color change o Endpoint: point where the indicator changes color ▪ Finding range of color change for an indicator • Lower range = pKa – 1 • Upper range = pKa + 1 ▪ You want to use an indicator whose pKa is closest to the equivalence point of the titration • Neutralization o Basis of neutralization: acids release H+ and bases release OH- which join to form water o Neutralizations are highly exothermic ▪ H+ + OH- → H2O is very product favored, K > 1 ▪ Standard free energy = -RTln(K) • If K>1, then G is negative • Interpretation of the titration curves o Graph is represented in a sigmoidal curve o Equivalence point: the midpoint of the titration curve ▪ For a monoprotic acid: when there's equal amounts of acid and base in solution ▪ When a strong acid is titrated with a strong base, the equivalence point is pH 7 ▪ For titrations where acid and base are of different strengths • If acid is weaker than the base, then the equivalence is >7 • If the acid is stronger than the base, then equivalence is <7 o Half-equivalence point: point where ½ of acid has been neutralized by the base (concentration of CB is equal to the concentration of remaining acid) ▪ Only happens for weak-strong acid-base titrations ▪ pH where the solution is most resistant to pH changes (best for buffering) ▪ pH of half-equivalence point = pKa of the weak acid • • o o Weak acid/weak-base titrations ▪ Similar shape to strong/weak acid/base titrations (half-eq and eq point) but compressed pH range Solving for pH's at points on the titration curve o Half-equivalence point ▪ Use of Henderson-Hasselbach equation (since pH of the solution at this point equals the pH of the weak acid) ▪ o Equivalence Point ▪ Find the pH by first finding the pOH and subtracting that from 14 ▪ Steps • 1) Find the Kb by dividing Kw and Ka • 2) Set Kb = to [OH-][HA]/[A-], solve to find the OH- concentration • 3) Find pOH, subtract that from 14 Redox titration o Strong oxidizing agent is titrated, resulting voltage change measured o Like a normal titration but measures changes in electrical potential, not pH, of a solution o Interpretation of curves ▪ Curve ascends slowly at first, then quite suddenly, then ascends slowly again ▪ Equivalence point: when voltage suddenly shoots up, indicates where all moles of a reducing agent have been oxidized