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16.S: Acid–Base Equilibria (Summary)

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    91292
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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 there is more H+ than OH-
      • Solution is basic if there is more OH- than 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

    16.4.1 Proton Transfer Reactions

      • Brq nsted-Lowry definition of acids an bases
      • Acid is a proton donor
      • Base is a proton acceptor
      • Can be applied to non-aqueous solutions
      • Brq nsted-Lowry acid must be able to lose a H+ ion
      • Brq nsted-Lowry base must have a non bonding pair of electrons to bind to H+ ion
      • Amphoteric - substance that can act as an acid or base

    16.4.2 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

    16.4.3 Related Strengths of Acids an 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
      • red litmus paper for pH of 5 or lower
      • blue litmus paper for pH of 8 or higher

    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) + H2O(l) ß à H3O+ + A-(aq)
      • HA(aq) ß à H+(aq) + A-(aq)
      • 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
      • 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 than 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 or 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 Brq nsted-Lowry is a Lewis base
      • Lewis acids have molecules that have incomplete octets

    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 4.0 license and was authored, remixed, and/or curated by LibreTexts.

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