Chapter 16 Acids & Bases

Acids vs. Bases Section 16.1 Arrhenius model Brønsted-Lowry model Ex: Oldest model; only applies to compounds that contain H+ or OH- ions Brønsted-Lowry model Refers to a compound’s ability to donate or accept an H+ ion Ex: HCl NH3 H2O

Brønsted-Lowry Acids and Bases Section 16.2 Water as Brønsted-Lowry acid/base: No such thing as H+ ion in solution (too unstable) Only H3O+ Proton transfer reactions:

Brønsted-Lowry Example NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Which acts as Brønsted-Lowry base? acid? H2O behaves as an amphoteric compound Capable of accepting OR donating H+ ions

Conjugate Acids/Bases The reactants and products associated with a proton transfer reaction are known as a conjugate acid-base pair:

Relative Strength of Acids and Bases The strength of an acid depends primarily on the willingness to donate or accept electrons i.e. strength of conjugate acid/base Equilibrium for strong acids lies heavily on the side of the deprotonated form and vice versa

The Autoionization of Water Section 16.3 Water has a very interesting property due to its amphoterism Capable of autoionization:

Calculating the [H3O+] An acid is added to water so that the hydrogen ion concentration is 0.25 M. Calculate the hydroxide ion concentration. See Sample Exercise 16.4 (Pg. 674)

The pH Scale Section 16.4 Concentrations of either H3O+ or OH- are typically very small and therefore cumbersome The pH scale is a logarithmic scale and is much more convenient: pH = -log[H3O+] pOH = -log[OH-] pH + pOH = 14

pH of Common Substances

Strong Acids and Bases Section 16.5 For a compound to be classified as a strong acid or base it must completely dissociate into ions when placed into aqueous solution Very weak conjugate bases No equilibrium Ex: HCl(g) + H2O(l)  H3O+(aq) + NO3-(aq) There are 7 strong acids which you will have to remember: HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4

HF(g) + H2O(l) H3O+(aq) + F-(aq) Weak Acids Section 16.6 Weak acids have a relatively strong conjugate base (compared to strong acids) and therefore some competition exists for the proton HF(g) + H2O(l) H3O+(aq) + F-(aq) As a result, the equilibrium between conjugate acid/base is treated as if it were any other equilibrium

Calculating Ka for a Weak Acid Picric acid, a weak acid, is dissolved in water to prepare a 0.100 M solution. The pH of the solution was found to be 1.09. Calculate Ka for picric acid at this temperature. See Sample Exercise 16.10 (Pg. 682)

Using Ka to Calculate pH Calculate the pH in a 0.100 M solution of the weak acid naphthol, for which Ka is 1.7  10-10 See Sample Exercise 16.12 (Pg. 685)

Using Ka to Calculate pH Calculate the pH of 0.017 M C6H5COOH. The Ka of C6H5COOH is 6.3 x 10-5. See Sample Exercise 16.12 (Pg. 685)

pH of Weak Acids Calculate the pH of a 0.10 M HF solution

Percent Ionization Weak acids, by definition, do not ionize 100% when placed in aqueous solution It is therefore possible to calculate the extent of ionization (percent ionization)

Calculating Percent Ionization Calculate the percent of HF molecules ionized in a 0.010 M HF solution. See Sample Exercise 16.13 (Pg. 687)

Effect of Concentration on Percent Ionization As the concentration of a weak acid increases, the equilibrium concentration of H3O+(aq) increases However, the percent ionization decreases as concentration increases

Polyprotic Acids Acids that have more than one acidic proton are referred to as polyprotic acids H2C2O4 + H2O HC2O4- + H3O+ HC2O4- + H2O C2O42- + H3O+ = 5.6 x 10-2 = 1.6 x 10-5

Calculating pH for Polyprotic Acid Solutions Because the acidic protons of a polyprotic acid are removed in discreet steps, we must calculate the [H3O+] for each step However, because Ka1 is typically very large compared to Ka2 or Ka3, we may initially assume that the [H3O+] is a result of the first ionization

Calculating pH and Concentration of All Species for a Polyprotic Acid Calculate the pH of a 0.0037 M H2CO3 solution. For super duper fun, calculate the [CO32-] in solution. See Sample Exercise 16.14 (Pg. 689)

Weak Bases Section 16.7 There are fewer weak bases compared to the number of weak acids Equilbria are the same (only involve OH- as product as opposed to H3O+)

Calculating pH of Weak Base Solutions What is the pH of a 1.44 M (concentration of household ammonia) solution of NH3? Kb = 1.8 x 10-5 See Sample Exercise 16.15 (Pg. 691)

Relationship Between Ka and Kb Section 16.8 Examine the following equilibria: NH4+(aq) NH3(aq) + H+(aq) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

Converting Between Ka and Kb The product of the dissociation constants for any conjugate acid/base pair is always equal to the dissociation constant for H2O: Ka x Kb = Kw = 1.0 x 10-14 Likewise: pKa + pKb = pKw = 14.00

Calculating pH of a Weak Base (Two Approaches) Calculate the pH of a solution that is 0.0500 M in ammonium ion. Kb = 1.8 x 10-5

Acid-Base Properties of Salt Solutions Section 16.9 Hydrolysis reactions The conjugate base of weak acids are capable of producing hydroxide ions in solution Raises pH Conjugate acid of weak bases are capable of producing hydronium ions in solution Lowers pH

Acid-Base Behavior and Chemical Structure Section 16.10 The chemical structure of a compound is what ultimately determines acid/base behavior Ex: Why does NaOH act as a base whereas CH3OH acts as a weak acid? Why does KH act as a strong base? Factors that affect? Polarization of H—X bond (NaH vs. CH4) Bond strength (HF vs. HCl) Stability of conjugate base (HNO3 vs. HNO2)

Oxyacids Special attention is paid to oxyacids because the majority of weak acids fall into this category The strength of these acids is governed by one overriding principle: charge stabilization Either through electronegativity or resonance

Effect of Additional Oxygens Because oxygen is a very electronegative element, the addition of more oxygen leads to a stronger acid:

Carboxylic Acids A second major class of weak acids are the carboxylic acids Characterized by the presence of an COOH group