Chemistry Honors Acids and Bases Lesson 6

Chemistry Honors Acids and Bases Definitions 1. Arrhenius Definition – The first definition of acids and bases. Arrhenius defined an acid as a species that produces H+ (a proton) in an aqueous solution and a base as a species that produces OH- (a hydroxide ion) in a aqueous solution. 2. Bronsted-Lowry Definition – A Bronsted-Lowry acid donates protons, while a Bronsted-Lowry base accepts protons. However, they can not be called Arrhenius bases since in aqueous solution they do not dissociate to form OH-. The advantage of this definition is that it is not limited to aqueous solutions. Bronsted-Lowry acids and bases always occur in pairs called conjugate acidbase pairs. The pairs are related through the transfer of a proton. 3. Lewis Definition – Lewis defined an acid as an electron-pair acceptor and a base as a electron pair donor. This definition is the most inclusive and encompasses species not included in the Bronsted-Lowry definition. Hydrogen Ion Equilibria Hydrogen ion concentration, [H+], is a measure of both Arrhenius and Bronsted-Lowry defined acids. This is generally expressed as pH, where: pH = -log[H+] = log(1/[H+]) Likewise, hydroxide ion concentration, [OH-], is measured as pOH. pOH = -log[OH-] = log(1/[OH-]) In any aqueous solution, the H2O solvent dissociates slightly: H2O <===> H+ + OHThis dissociation is an equilibrium reaction and is therefore described by a constant, Kw, the water dissociation constant. Kw = [H+][OH-] = 10-14 Rewriting this equation using logarithms show: pH + pOH = 14 In pure H20, [H+] is equal to [OH-], since for everyone H20 that dissociates, one H+ and OH- is made. A solution with equal H+ and OH- is neutral and has a pH of 7. A pH below 7 indicates an excess of H+ and thus an acidic solution. A pH above 7 indicates a relative excess of OH- ions and therefore a basic solution. Strong Acids and Bases Strong acids and bases are those that completely dissociate into their component ions in aqueous solution. For example, NaOH is added to water, it dissociates completely: NaOH + water ---> Na+ + OHIn a 1 M solution of NaOH, complete dissociation gives 1 M of OH-: PH = 14 – (-log[OH-]) = 14 + log[1] = 14 Weak Acids and Bases Weak acids and bases are those that only partially dissociate in aqueous solution. A weak monoprotic acid, HA in aqueous solution will achieve the follwing equilibrium after dissociation: HA + H2O <===> H3O+ + AThe acid dissociation constant, Ka, is a measure of the how much the acid dissociates. Ka = [H30+][A-]/[HA] Similarly, a weak monovalent base, BOH, dissociates to give B+ and OH-. The base dissociation constant, Kb, is a measure of the degree that the base dissociates. Kb = [B+][OH-]/[BOH] A conjugate acid is define as the acid formed when a base gains a proton. Similarly, a conjugate base is formed when an acid loses a proton. For example, with HCO3-/CO32conjugate acid/base pair, the CO32- is the conjugate base and HCO3- is the conjugate acid. By adding the reactions of the conjugate acid and conjugate base with water, the net reaction is simply the dissociation of water. Thus if you know the dissociation constant for one, the other can be determined: Ka x Kb = Kw = 1x10-14 Thus Ka and Kb are inversely related. In other words, if Ka is large (the acid is trong), then Kb (base is weak) will be small and vice versa. From this relationship, one can see that when a conjugate acid/base pair form from a weak acid, the conjugate base is stronger than the acid. Titration Titration is a procedure used to determine the concentration of an acid or base. This is accomplished by reacting a known volume of a solution of an unknown concentration with a known volume of a solution of known concentration. When the number of acid molecules equal the number of base molecules added (or vice versa), the equivalence point is reached. The equivalence point in a titration is estimated in two common ways.: either by using a graphical method, plotting the pH of the solution as a function of added titrant by using a pH meter or by watching for a color change of an added indicator. Indicators are weak organic acids or bases that have different colors in their undissociated and dissociated states. These states are dependent on pH. The Henderson-Hasselbach equation is used to estimate the pH of a solution: PH = pKa + log[conjugate base]/[weak acid] Note that when [conjugate base] = [weak acid] the pH = pKa because the log 1 = 0. In a titration, half-way to the equivalent point is when [conjugate base] = [weak acid].