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6.S: Acid–Base Equilibrium (Study Guide)

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    6.1: Brønsted–Lowry Acids and Bases

    • Brønsted-Lowry definition of acids an bases
    • Acid is a proton donor
    • Base is a proton acceptor
    • Can be applied to non-aqueous solutions
    • Brønsted-Lowry acid must be able to lose a H+ ion
    • Brønsted-Lowry base must have at least one non-bonding pair (lone pair) of electrons to bind to H+ ion
    • Amphoteric - substance that can act as an acid or base
    • Conjugate Acid-Base Pairs
    • conjugate acid - product formed by adding a proton to base
    • conjugate base - product formed by removal of a proton from acid
    • autoionization of water - dissociation of H2O molecules to H+ and OH- ions
    • at room temperature only 1 out of 109 molecules are ionized
    • exclude water from equilibrium expressions involving aqueous solutions
    • ion-product constant
    • kw = k[H2O] = [H+][OH-] = 1.0 x 10-14 (at 25° C)
    • solution is neutral when [H+] = [OH-]
    • solution is acidic when [H+] > [OH-]
    • solution is basic when [H+] < [OH-]

    6.2: pH and pOH

    • concentration of [H+] expressed in terms of pH
    • pH = -log [H+]
    • acidic solutions [H+] > 1.0 x 10-7 [OH-] < 1.0 x 10-7 pH < 7.00
    • neutral solutions [H+] = [OH-] = 1.0 x 10-7 pH = 7
    • basic solutions [H+] < 1.0 x 10-7 [OH-] > 1.0 x 10-7 pH > 7

    Other "p" Series

    • pOH = -log [OH-]
    • pH + pOH = -log Kw = 14.00

    6.3 Relative Strengths of Acids and Bases

    • the stronger the acid, the weaker the conjugate base
    • the stronger the base, the weaker the conjugate acid
    • equilibrium favors transfer of proton from stronger acid to stronger base

    Weak Acids

    • \(HA_{(aq)} + H_2O_{(l)} \to H_3O^+ + A^-_{(aq)}\)
    • \(HA_{(aq)} \to H^+_{(aq)} + A^-_{(aq)}\)
    • \(K_a = \frac{[H^+][A^-]}{[HA]}\)
    • Ka = acid - dissociation constant
    • The lager the Ka the stronger the acid
    • Ka usually less than 10-3

    Weak Bases

    • base-dissociation constant, Kb
    • equilibrium at which base reacts with H2O to form a conjugate acid and OH-
    • contain 1 or more lone pair of electrons

    Relationship Between Ka and Kb

    • when two reactions are added together then equilibrium constant of third reaction is equal to the product of the equilibrium constants of the added reactions
    • reaction 1 + reaction 2 = reaction 3
    • K1 x K2 = K3
    • Ka x Kb = [H+][OH-] = Kw
    • Acid-dissociation constant times base-dissociation constant equals the ion-product constant for water
    • Ka x Kb = Kw = 1.0 x 10-14
    • pKa x pKb = pKw = 14; (pKa= -log Ka and pKb = -log Kb)

    Calculating Ka or Kb given initial and equilibrium concentration

    1.Identify species, write dissociation equation

    2.Obtain K expression

    3.If pH or pOH given, calculate [H3O+] or [OH-]

    4.Build the corresponding ICE table

    5.Solve for all concentrations, answer question

    Calculating pH for Solution of Weak Acid or Base

    1) write ionization equilibrium

    2) write equilibrium expression

    3) I.C.E. Table

    4) substitute equilibrium concentrations into equilibrium expression

    • percent ionization = fraction of weak acid molecules that ionize * 100%
    • in weak acids [H+] is small fraction of concentration of acid (can make assumptions)
    • percent ionization depends on temperature, identity of acid and concentration
    • as percent ionization decreases, concentration increases

    6.4 Polyprotic Acids

    • more than one ionizable H atom
    • easier to remove first proton than second
    • acid dissociation constants are Ka1, Ka2, etc…
    • Ka values usually differ by 103

    6.5 Acid-Base Properties of Salt Solutions

    • hydrolysis - ions reacting with water to produce H+ and OH- ions
    • anions from weak acids react with water to produce OH- ions which is basic
    • anions of strong acids are not basic and do not influence pH
    • anions that have ionizable protons are amphoteric
    • behavior depends on Ka and Kb
    • all cations except those of alkali metals and heavier alkaline earth (Ca2+, Sr2+ and Ba2+) are weak acids in water
    • alkali metal and alkaline earth cations do not hydrolyze
    • do not affect pH
    • strengths of acids and bases from salts
    • 1) salts derived from strong acid and base
      • no hydrolysis and solution has pH of 7
    • 2) salts derived from strong base and weak acid
      • strong conjugate base
      • anion hydrolyzes and produces OH- ions
      • cation does not hydrolyze
      • pH greater than 7
    • 3) salts derived from weak base and strong acids
      • cation is strong conjugate acid
      • cation hydrolyzes to produce H+
      • anion does not hydrolyze
      • solution has pH below 7
    • 4) salts derived from weak acid and base
      • both cation and anion hydrolyze
      • pH depends on extent on hydrolysis of each ion]

    6.6 Lewis Acids and Bases

    • Lewis acid - electron pair acceptor
    • Lewis base - electron pair donor
    • Any Bronsted-Lowry base is a Lewis base
    • Lewis acids contain at least one atom with an incomplete octet

    6.S: Acid–Base Equilibrium (Study Guide) is shared under a CC BY-NC-SA 3.0 license and was authored, remixed, and/or curated by LibreTexts.