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16.S: Acid–Base Equilibria (Summary)

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     16.1: Acids and Bases: A Brief Review

    • acids have sour taste and turn litmus paper red
    • bases have a bitter taste and feel slippery
    • Svante Arrhenius (1859-1927)
    • Acids associated with H+ ions
    • Bases associated with OH- ions
    • Solution is acidic if [H+] > [OH-]
    • Solution is basic if [OH-] > [H+]

    16.2: Brønsted–Lowry Acids and Bases

    • Arrhenius definition of acids and bases
      • Acids when dissolved in water increase H+ concentration
      • Bases when dissolved in water increase OH- concentration

    16.2.1 Proton Transfer Reactions

    • 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

    16.2.2 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

    16.2.3 Related 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

    16.3: The Autoionization of Water

    • 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-]

    16.3.1 The Proton in Water

    • H+ ion is a proton with no valence electrons
    • H+ ion react with H2O molecule to form H3O+, hydronium ion
    • H3O+ ion can bond with other H2O molecules to form hydrated hydrogen ions
    • H+ and H3O+ used interchangeably

    16.4: The pH Scale

    • 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

    Measuring pH

    • pH meter
      • has a pair of electrodes connected to a meter that measures in millivolts
      • voltage generated when electrodes placed in solution, and is measured by meter
    • blue litmus paper turns red in acidic solution
    • red litmus paper turns blue in basic solution

    16.5: Strong Acids and Bases

    strong acids and bases are strong electrolytes

    16.5.1 Strong Acids

    strongest monoprotic acids

    • HCl, HBr, HI, HNO3, HclO3, HclO4, and diprotic H2SO4
    • For strong monoprotic acid concentration of [H+] equals the original concentration of the acid

    16.5.2 Strong Bases

    • most common strong bases are ionic hydroxides of alkali metals and the heavier alkaline-earth metals
    • complete dissociation

    16.6: 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

    16.6.1 Calculating pH for Solutions of Weak Acids

    • 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
      • percent ionization depends on temperature, identity of acid and concentration
      • as percent ionization decreases, concentration increases

    16.6.2 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

    16.7: 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

    16.7.1 Types of Weak Bases

    • weak bases have NH3 and anions of weak acids

    16.8: 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)

    16.9: 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]

    16.10: Acid-Base Behavior and Chemical Structure

    16.10.1 Factors that Affect Acid Strength

    • strength of acid depends on:
      • 1) polarity of H-X bond
      • 2) strength of H-X bond
      • 3) stability of conjugate base, X-
    • molecule will transfer proton if H-X bond is polarized
    • in ionic hydrides H- acts as proton acceptor because of negative charge
    • nonpolar bonds produce neither acidic nor basic solutions
    • strong bonds less easily dissociated that weak bonds
    • the greater the stability of conjugate base, the stronger the acid]

    16.10.2 Binary Hydrides

    • metal hydrides are basic or have no acid-base properties in water
    • nonmetal hydrides can be between having no acid-base properties to being acidic
    • in each group of nonmetallic elements, acidity increases with increasing atomic number
      • bond strengths decrease as central atom gets larger and overlap of orbitals get smaller

    16.10.3 Oxyacids

    • Y-O-H bond
    • Oxyacids - have OH bonded to central atom
    • Base if bonded to a metal because pair of electrons shared between Y-O is completely transferred to O
      • Ionic compound with OH- is formed
    • When bonded to nonmetal the bond is covalent and compounds are acidic or neutral
    • As electronegativity of Y increases , acidity also increases
      • O-H bond becomes more polar
      • Conjugate base usually an anion and stability increases as electronegativity of Y increases
    • Relating acid strengths of oxyacids to electronegativity of Y and to number of groups attached to Y
      • 1) same number of oxygen atoms, acid strength increases as electronegativity of central atom increases
      • 2) same central atom Y, acid strength increases with increasing number of bonded oxygen atoms to central atom
        • acidity increases as oxidation number of central atom increases

    16.10.4 Carboxylic Acids

    • carboxyl group - COOH
    • acidic behavior of carboxylic acids
      • addition oxygen atom in carboxyl group draws density from O-H bond which increases the polarity
      • conjugate base ion have resonance forms
      • acidity increases as number of electronegative atoms in acid increases

    16.11: 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

    16.11.1 Hydrolysis of Metal Ions

    • hydration - attraction of metal ions to water molecules
      • metal ion acts as Lewis acid
      • water molecule acts as Lewis base
      • electron density drawn from oxygen atom to water molecule
      • O-H bond becomes more polarized
    • For hydrolysis reactions Ka increases with increasing charge and decreasing radius of ion

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

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