3.091 Recitation 23: Acids and Bases

3.091: Introduction to Solid State Chemistry, Fall 2018 Recitation 23 1 Acids and bases In lecture, we talked about Brønsted-Lowry acids and bases as well as Lewis acids and bases. The BrønstedLowry definition of acids and bases is more narrow: a Brønsted-Lowry acid is a proton donor, while a Brønsted-Lowry base is a proton acceptor : to be clear here, the proton is a hydrogen ion, H+ . More generally, a Lewis acid is an electron acceptor, while a Lewis base is an electron donor. Water is amphoteric: it can be either an acid or a base. Water serves as the solvent for many acid-base reactions. Consider adding a generic acid to water: HA(aq) + H2 O(l) → H3 O+ (aq) + A− (aq) Similarly, consider adding a generic base to water: B − (aq) + H2 O(l) → HB(aq) + OH − (aq) These serve as prototypical acid-base reactions: we can identify conjugate acid-base pairs to relate species that lost/gained a proton (or electron) on either side of the reaction. Example: Identify the conjugate acid-base pairs in the following reaction: CH3 CO2 H(aq) + N H3 (aq) → CH3 CO2− + N H4+ 2 Dissociation In the same way we wrote a general equilibrium constant, we can write an acid/base specific equivalent. The equilibrium constant for our acid dissociation reaction above would be Keq = [H3 O+ ][A− ] [H2 O][HA] However, since the water is acting as a solvent and is present in excess, it has a constant concentration: we therefore define the acid constant as Ka = Keq [H2 O] = [H3 O+ ][A− ] [HA] Similarly, the base dissocation constant is Kb = [HB][OH − ] [B−] 1 3.091: Introduction to Solid State Chemistry, Fall 2018 Recitation 23 The product of the acid and base constants for a given dissociation is constant and called KW , where W is for water. KW = 10−14 The stronger the acid or base, the more readily it dissolves in solution, and the greater the magnitude of Ka or Kb . There are only a few strong acids– acids that fully dissociate into their constituent ions. These include HCl, HBr, HI, HNO3 , HClO3 , HCLO4 , and H2 SO4 . The strong bases are mostly alkali salts, including LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2 , Sr(OH)2 , Ba(OH)2 . Of course, the strength of an acid or base is related to its bond energy: polar acids with big differences in electronegativity have a low bond energy rapidly dissociate when placed in water, for example. Example: Write an expression for the dissociation and for the acid or base dissociation constants for these acids (HF, CH3 CO2 H) and bases (NH3 , NaOH). Hydrofluoric acid, HF, is a strong acid. In water, the proton dissociates: HF (aq) + H2 O(l) → H3 O+ (aq) + F − (aq) Its acid dissociation constant is Ka = [H3 O+ ][F − ] [HF ] We can take a similar approach for acetic acid, a weak acid. CH3 CO2 H(aq) + H2 O(l) → H3 O+ (aq) + CH3 CO2− (aq) The acid dissociation constant is Ka = [H3 O+ ][CH3 CO2− ] [CH3 CO2 H] For ammonia, the base gains a proton when placed in water: N H3 (aq) + H2 O(l) → N H4+ (aq) + OH − (aq) The base dissociation constant for ammonia is Kb = [N H4+ ][OH − ] [N H3 ] Finally, sodium hydroxide fully dissociates: N aOH(aq) + H2 O(l) → N a+ (aq) + OH − (aq) The base dissociation constant is Kb = [N a+][OH − ] [N aOH] Example: Which of these acids is the strongest? sulfurous acid (H2 SO3 ; Ka = 1.54×10−2 ), phosphoric −8 acid (H2 PO− ), citric acid (H3 C6 H5 O7 ; Ka = 8.4×10−4 ) 4 ; Ka =6.23×10 Stronger acids have larger acid dissociation constants, so sulfurous acid is stronger than citric acid and phosphoric acid. 2