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Chapter 10


Acids, Bases,
  and Salts
Chapter 10
      Table of Contents
      10.1Arrhenius Acid-Base Theory
      10.2Brønsted-Lowry Acid-Base Theory
      10.3Mono-, Di-, and Triprotic Acids
      10.4Strengths of Acids and Bases
      10.5Ionization Constants for Acids and Bases
      10.6Salts
      10.7Acid-Base Neutralization Reactions
      10.8Self-Ionization of Water
      10.9The pH Concept
      10.10The pKa Method for Expressing Acid Strength
      10.11The pH of Aqueous Salt Solutions
      10.12Buffers
      10.13The Henderson-Hasselbalch Equation
      10.14Electrolytes
      10.15Equivalents and Milliequivalents of Electrolytes
      10.16Acid-Base Titrations
Copyright © Cengage Learning. All rights reserved             2
Section 10.1
      Arrhenius Acid-Base Theory



      • Arrhenius acid: hydrogen-containing compound
        that produces H+ ions in solution.
          Example: HNO3 → H+ + NO3–
      • Arrhenius base: hydroxide-containing compound
        that produces OH– ions in solution.
          Example: NaOH → Na+ + OH–




                                                        Return to TOC


Copyright © Cengage Learning. All rights reserved               3
Section 10.1
      Arrhenius Acid-Base Theory

      Ionization
      • The process in which individual positive and
        negative ions are produced from a molecular
        compound that is dissolved in solution.
         – Arrhenius acids




                                                       Return to TOC


Copyright © Cengage Learning. All rights reserved              4
Section 10.1
      Arrhenius Acid-Base Theory

      Dissociation
      • The process in which individual positive and
        negative ions are released from an ionic
        compound that is dissolved in solution.
         – Arrhenius Bases




                                                       Return to TOC


Copyright © Cengage Learning. All rights reserved              5
Section 10.1
      Arrhenius Acid-Base Theory

      Difference Between Ionization and Dissociation




                                                       Return to TOC


Copyright © Cengage Learning. All rights reserved              6
Section 10.2
      Brønsted-Lowry Acid-Base Theory



      • Brønsted-Lowry acid: substance that can donate
        a proton (H+ ion) to some other substance;
        proton donor.
      • Brønsted-Lowry base: substance that can
        accept a proton (H+ ion) from some other
        substance; proton acceptor.
                   HCl + H2O → Cl− + H3O+
                                       acid         base


                                                           Return to TOC


Copyright © Cengage Learning. All rights reserved                  7
Section 10.2
      Brønsted-Lowry Acid-Base Theory

      Brønsted-Lowry Reaction




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved           8
Section 10.2
      Brønsted-Lowry Acid-Base Theory

      Acid in Water



        HA(aq) +                           H2O(l)   H3O+(aq) +    A-(aq)
           acid                              base   conjugate    conjugate
                                                            acid
              base




                                                                           Return to TOC


Copyright © Cengage Learning. All rights reserved                                  9
Section 10.2
      Brønsted-Lowry Acid-Base Theory

      Acid Ionization Equilibrium




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          10
Section 10.2
      Brønsted-Lowry Acid-Base Theory

      Amphiprotic Substance
      • A substance that can either lose or accept a
        proton and thus can function as either a
        Brønsted-Lowry acid or a Brønsted-Lowry base.
          Example: H2O, H3O+
                                      H2O, OH–




                                                        Return to TOC


Copyright © Cengage Learning. All rights reserved              11
Section 10.3
      Mono-, Di-, and Triprotic Acids

      Monoprotic Acid
      • An acid that supplies one proton (H+ ion) per
        molecule during an acid-base reaction.

                     HA + H2O                       A− + H3O+




                                                                Return to TOC


Copyright © Cengage Learning. All rights reserved                      12
Section 10.3
      Mono-, Di-, and Triprotic Acids

      Diprotic Acid
      • An acid that supplies two protons (H+ ions) per
        molecule during an acid-base reaction.

                    H2A + H2O                       HA− + H3O+
                    HA− + H2O                       A2− + H3O+




                                                                 Return to TOC


Copyright © Cengage Learning. All rights reserved                       13
Section 10.3
      Mono-, Di-, and Triprotic Acids

      Triprotic Acid
      • An acid that supplies three protons (H+ ions) per
        molecule during an acid-base reaction.

                    H3A + H2O                       H2A− + H3O+
                    H2A− + H2O                      HA2− + H3O+
                    HA2− + H2O                      A3− + H3O+




                                                                  Return to TOC


Copyright © Cengage Learning. All rights reserved                        14
Section 10.3
      Mono-, Di-, and Triprotic Acids

      Polyprotic Acid
      • An acid that supplies two or more protons (H+
        ions) during an acid-base reaction.
      • Includes both diprotic and triprotic acids.




                                                        Return to TOC


Copyright © Cengage Learning. All rights reserved              15
Section 10.4
      Strengths of Acids and Bases

      Strong Acid
      • Transfers ~100% of its protons to water in an
        aqueous solution.
      • Ionization equilibrium lies far to the right.
      • Yields a weak conjugate base.




                                                        Return to TOC


Copyright © Cengage Learning. All rights reserved              16
Section 10.4
      Strengths of Acids and Bases

      Commonly Encountered Strong Acids




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          17
Section 10.4
      Strengths of Acids and Bases

      Weak Acid
      • Transfers only a small % of its protons to water
        in an aqueous solution.
      • Ionization equilibrium lies far to the left.
      • Weaker the acid, stronger its conjugate base.




                                                           Return to TOC


Copyright © Cengage Learning. All rights reserved                 18
Section 10.4
      Strengths of Acids and Bases
      Differences Between Strong and Weak Acids in Terms of Species
      Present




                                                                      Return to TOC


Copyright © Cengage Learning. All rights reserved                            19
Section 10.4
      Strengths of Acids and Bases

      Bases
      • Strong bases: hydroxides of Groups IA and IIA.




                                                         Return to TOC


Copyright © Cengage Learning. All rights reserved               20
Section 10.5
      Ionization Constants for Acids and Bases

      Acid Ionization Constant
      • The equilibrium constant for the reaction of a
        weak acid with water.

                      HA(aq) + H2O(l)                           H3O+(aq) + A-(aq)

                                                         H3O+   A − 
                                                    Ka =       
                                                            [ HA ]


                                                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved                                          21
Section 10.5
      Ionization Constants for Acids and Bases

      Acid Strength, % Ionization, and Ka Magnitude

      • Acid strength increases as % ionization
        increases.
      • Acid strength increases as the magnitude of Ka
        increases.
      • % ionization increases as the magnitude of Ka
        increases.




                                                         Return to TOC


Copyright © Cengage Learning. All rights reserved               22
Section 10.5
      Ionization Constants for Acids and Bases

      Base Ionization Constant
      • The equilibrium constant for the reaction of a
        weak base with water.

                      B(aq) + H2O(l)                         BH+(aq) + OH–(aq)

                                                         BH+  OH− 
                                                    Kb =          
                                                             [ B]


                                                                                 Return to TOC


Copyright © Cengage Learning. All rights reserved                                       23
Section 10.6
      Salts



      • Ionic compounds containing a metal or
        polyatomic ion as the positive ion and a
        nonmetal or polyatomic ion (except hydroxide)
        as the negative ion.
      • All common soluble salts are completely
        dissociated into ions in solution.




                                                        Return to TOC


Copyright © Cengage Learning. All rights reserved              24
Section 10.7
      Acid-Base Neutralization Reactions

      Neutralization Reaction
      • The chemical reaction between an acid and a
        hydroxide base in which a salt and water are the
        products.

                                      HCl + NaOH → NaCl + H2O


                             H2SO4 + 2 KOH → K2SO4 + 2 H2O


                                                                Return to TOC


Copyright © Cengage Learning. All rights reserved                      25
Section 10.7
      Acid-Base Neutralization Reactions

      Formation of Water




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          26
Section 10.8
      Self-Ionization of Water

      Self-Ionization
      • Water molecules in pure water interact with one
        another to form ions.

                                      H2O + H2O     H3O+ + OH–


      • Net effect is the formation of equal amounts of
        hydronium and hydroxide ions.


                                                                 Return to TOC


Copyright © Cengage Learning. All rights reserved                       27
Section 10.8
      Self-Ionization of Water

      Self-Ionization of Water




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          28
Section 10.8
      Self-Ionization of Water

      Ion Product Constant for Water
      • At 24°C:
            Kw = [H3O+][OH–] = 1.00 × 10–14
      • No matter what the solution contains, the
        product of [H3O+] and [OH–] must always equal
        1.00 × 10–14.




                                                        Return to TOC


Copyright © Cengage Learning. All rights reserved              29
Section 10.8
      Self-Ionization of Water

      Relationship Between [H3O+] and [OH–]




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          30
Section 10.8
      Self-Ionization of Water

      Three Possible Situations

      • [H3O+] = [OH–]; neutral solution
      • [H3O+] > [OH–]; acidic solution
      • [H3O+] < [OH–]; basic solution




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          31
Section 10.8
      Self-Ionization of Water

                        Exercise


           Calculate [H3O+] or [OH–] as required for each
           of the following solutions at 24°C, and state
           whether the solution is neutral, acidic, or basic.

             b) 1.0 × 10–4 M OH–
                      1.0 × 10–10 M H3O+; basic
             b) 2.0 M H3O+
                     5.0 × 10–15 M OH–; acidic
                                                                Return to TOC


Copyright © Cengage Learning. All rights reserved                      32
Section 10.9
      The pH Concept



      • pH = –log[H3O+]
      • A compact way to represent solution acidity.
      • pH decreases as [H+] increases.
      • pH range between 0 to 14 in aqueous solutions
        at 24°C.




                                                        Return to TOC


Copyright © Cengage Learning. All rights reserved              33
Section 10.9
      The pH Concept

                        Exercise


     Calculate the pH for each of the following
      solutions.

             a) 1.0 × 10–4 M H3O+
                      pH = 4.00
             – 0.040 M OH–
                     pH = 12.60


                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          34
Section 10.9
      The pH Concept

                        Exercise


           The pH of a solution is 5.85. What is the [H3O+]
           for this solution?

                    [H3O+] = 1.4 × 10–6 M




                                                              Return to TOC


Copyright © Cengage Learning. All rights reserved                    35
Section 10.9
      The pH Concept

      pH Range
      • pH = 7; neutral
      • pH > 7; basic
         – Higher the pH, more basic.
      • pH < 7; acidic
         – Lower the pH, more acidic.




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          36
Section 10.9
      The pH Concept

      Relationships Among pH Values, [H3O+], and [OH–]




                                                         Return to TOC


Copyright © Cengage Learning. All rights reserved               37
Section 10.10
      The pKa Method for Expressing Acid Strength



      • pKa = –log Ka
      • pKa is calculated from Ka in exactly the same
        way that pH is calculated from [H3O+].




                                                        Return to TOC


Copyright © Cengage Learning. All rights reserved              38
Section 10.10
      The pKa Method for Expressing Acid Strength

                        Exercise


                  Calculate the pKa for HF given that the Ka for
                  this acid is 6.8 × 10–4.

                                                    pKa = 3.17




                                                                   Return to TOC


Copyright © Cengage Learning. All rights reserved                         39
Section 10.11
      The pH of Aqueous Salt Solutions

      Salts
      • Ionic compounds.
      • When dissolved in water, break up into its ions
        (which can behave as acids or bases).
      • Hydrolysis – the reaction of a salt with water to
        produce hydronium ion or hydroxide ion or both.




                                                            Return to TOC


Copyright © Cengage Learning. All rights reserved                  40
Section 10.11
      The pH of Aqueous Salt Solutions

      Types of Salt Hydrolysis
      • The salt of a strong acid and a strong base does
        not hydrolyze, so the solution is neutral.
          KCl, NaNO3




                                                           Return to TOC


Copyright © Cengage Learning. All rights reserved                 41
Section 10.11
      The pH of Aqueous Salt Solutions

      Types of Salt Hydrolysis
      • The salt of a strong acid and a weak base
        hydrolyzes to produce an acidic solution.
          NH4Cl


                    NH4+ + H2O → NH3 + H3O+




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          42
Section 10.11
      The pH of Aqueous Salt Solutions

      Types of Salt Hydrolysis
      • The salt of a weak acid and a strong base
        hydrolyzes to produce a basic solution.
          NaF, KC2H3O2


                    F– + H2O → HF + OH–


                    C2H3O2– + H2O → HC2H3O2 + OH–


                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          43
Section 10.11
      The pH of Aqueous Salt Solutions

      Types of Salt Hydrolysis
      • The salt of a weak acid and a weak base
        hydrolyzes to produce a slightly acidic, neutral,
        or slightly basic solution, depending on the
        relative weaknesses of the acid and base.




                                                            Return to TOC


Copyright © Cengage Learning. All rights reserved                  44
Section 10.11
      The pH of Aqueous Salt Solutions

      Neutralization “Parentage” of Salts




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          45
Section 10.12
      Buffers

      Key Points about Buffers
      • Buffer – an aqueous solution containing
        substances that prevent major changes in
        solution pH when small amounts of acid or base
        are added to it.
      • They are weak acids or bases containing a
        common ion.
      • Typically, a buffer system is composed of a weak
        acid and its conjugate base.


                                                       Return to TOC


Copyright © Cengage Learning. All rights reserved             46
Section 10.12
      Buffers

      Buffers Contain Two Active Chemical Species

      1. A substance to react with and remove added
         base.
      2. A substance to react with and remove added
         acid.




                                                      Return to TOC


Copyright © Cengage Learning. All rights reserved            47
Section 10.12
      Buffers

      Adding an Acid to a Buffer




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          48
Section 10.12
      Buffers

      Buffers




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          49
Section 10.12
      Buffers

      Addition of Base [OH– ion] to the Buffer

                                      HA + H2O      H3O+ + A–
      • The added OH– ion reacts with H3O+ ion,
        producing water (neutralization).
      • The neutralization reaction produces the stress
        of not enough H3O+ ion because H3O+ ion was
        consumed in the neutralization.
      • The equilibrium shifts to the right to produce
        more H3O+ ion, which maintains the pH close to
        its original level.
                                                                Return to TOC


Copyright © Cengage Learning. All rights reserved                      50
Section 10.12
      Buffers

      Addition of Acid [H3O+ ion] to the Buffer

                                      HA + H2O      H3O+ + A–
      • The added H3O+ ion increases the overall
        amount of H3O+ ion present.
      • The stress on the system is too much H3O+ ion.
      • The equilibrium shifts to the left consuming most
        of the excess H3O+ ion and resulting in a pH
        close to the original level.


                                                                Return to TOC


Copyright © Cengage Learning. All rights reserved                      51
Section 10.13
      The Henderson-Hasselbalch Equation

      Henderson-Hasselbalch Equation


                                                               −
                                                            A 
                                            pH = pK a + log  
                                                            [ HA ]




                                                                     Return to TOC


Copyright © Cengage Learning. All rights reserved                           52
Section 10.13
      The Henderson-Hasselbalch Equation

                        Exercise


           What is the pH of a buffer solution that is 0.45
           M acetic acid (HC2H3O2) and 0.85 M sodium
           acetate (NaC2H3O2)? The Ka for acetic acid is
           1.8 × 10–5.

                                                    pH = 5.02




                                                                Return to TOC


Copyright © Cengage Learning. All rights reserved                      53
Section 10.14
      Electrolytes


      • Acids, bases, and soluble salts all produce ions
        in solution, thus they all produce solutions that
        conduct electricity.
      • Electrolyte – substance whose aqueous solution
        conducts electricity.




                                                            Return to TOC


Copyright © Cengage Learning. All rights reserved                  54
Section 10.14
      Electrolytes

      Nonelectrolyte – does not conduct electricity
      • Example: table sugar (sucrose), glucose




                                                      Return to TOC


Copyright © Cengage Learning. All rights reserved            55
Section 10.14
      Electrolytes

      Strong Electrolyte – completely ionizes/dissociates
      • Example: strong acids, bases, and soluble salts




                                                            Return to TOC


Copyright © Cengage Learning. All rights reserved                  56
Section 10.14
      Electrolytes

      Weak Electrolyte – incompletely ionizes/dissociates
      • Example: weak acids and bases




                                                            Return to TOC


Copyright © Cengage Learning. All rights reserved                  57
Section 10.15
      Equivalents and Milliequivalents of Electrolytes

      Equivalent (Eq) of an Ion
      • The molar amount of that ion needed to supply
        one mole of positive or negative charge.

                                      1 mole K+ = 1 equivalent
                                      1 mole Mg2+ = 2 equivalents
                                      1 mole PO43– = 3 equivalents




                                                                     Return to TOC


Copyright © Cengage Learning. All rights reserved                           58
Section 10.15
      Equivalents and Milliequivalents of Electrolytes

      Milliequivalent


                      1 milliequivalent = 10–3 equivalent




                                                            Return to TOC


Copyright © Cengage Learning. All rights reserved                  59
Section 10.15
      Equivalents and Milliequivalents of Electrolytes

      Concentrations of Major Electrolytes in Blood Plasma




                                                             Return to TOC


Copyright © Cengage Learning. All rights reserved                   60
Section 10.15
      Equivalents and Milliequivalents of Electrolytes

                        Exercise


                   The concentration of Ca2+ ion present in a
                   sample is 5.3 mEq/L. How many milligrams
                   of Ca2+ ion are present in 180.0 mL of the
                   sample?

                                                     19 mg Ca2+ ion


               (             )(              )(                )(                   )(                     )(             )
                                                                               2+                     2+

 ( 180 mL )          1L
                   1000 mL
                                  5.3 mEq
                                     1L
                                                      1Eq
                                                    1000 mEq
                                                                    1 mol Ca
                                                                    2 Eq Ca
                                                                              2+
                                                                                         40.08 g Ca
                                                                                         1 mol Ca
                                                                                                    2+
                                                                                                                1000 mg
                                                                                                                  1g
                                                                                                                              = 19 mg Ca
                                                                                                                                           2+
                                                                                                                                                 ion



                                                                                                                                                Return to TOC


Copyright © Cengage Learning. All rights reserved                                                                                                      61
Section 10.16
      Acid-Base Titrations


      • A neutralization reaction in which a measured
        volume of an acid or a base of known
        concentration is completely reacted with a
        measured volume of a base or an acid of
        unknown concentration.
      • For a strong acid and base reaction:
           H+(aq) + OH–(aq) → H2O(l)



                                                        Return to TOC


Copyright © Cengage Learning. All rights reserved              62
Section 10.16
      Acid-Base Titrations

      Titration Setup




                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved          63
Section 10.16
      Acid-Base Titrations

      Acid-Base Indicator
      • A compound that exhibits different colors
        depending on the pH of its solution.
      • An indicator is selected that changes color at
        a pH that corresponds as nearly as possible
        to the pH of the solution when the titration is
        complete.




                                                          Return to TOC


Copyright © Cengage Learning. All rights reserved                64
Section 10.16
      Acid-Base Titrations

      Indicator – yellow in acidic solution; red in basic solution




                                                                     Return to TOC


Copyright © Cengage Learning. All rights reserved                           65
Section 10.16
      Acid-Base Titrations

                        Concept Check


        For the titration of sulfuric acid (H2SO4) with
        sodium hydroxide (NaOH), how many moles of
        sodium hydroxide would be required to react
        with 1.00 L of 0.500 M sulfuric acid to reach the
        endpoint?

                                                    1.00 mol NaOH


                                                                    Return to TOC


Copyright © Cengage Learning. All rights reserved                          66

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Chapter10

  • 2. Chapter 10 Table of Contents 10.1Arrhenius Acid-Base Theory 10.2Brønsted-Lowry Acid-Base Theory 10.3Mono-, Di-, and Triprotic Acids 10.4Strengths of Acids and Bases 10.5Ionization Constants for Acids and Bases 10.6Salts 10.7Acid-Base Neutralization Reactions 10.8Self-Ionization of Water 10.9The pH Concept 10.10The pKa Method for Expressing Acid Strength 10.11The pH of Aqueous Salt Solutions 10.12Buffers 10.13The Henderson-Hasselbalch Equation 10.14Electrolytes 10.15Equivalents and Milliequivalents of Electrolytes 10.16Acid-Base Titrations Copyright © Cengage Learning. All rights reserved 2
  • 3. Section 10.1 Arrhenius Acid-Base Theory • Arrhenius acid: hydrogen-containing compound that produces H+ ions in solution.  Example: HNO3 → H+ + NO3– • Arrhenius base: hydroxide-containing compound that produces OH– ions in solution.  Example: NaOH → Na+ + OH– Return to TOC Copyright © Cengage Learning. All rights reserved 3
  • 4. Section 10.1 Arrhenius Acid-Base Theory Ionization • The process in which individual positive and negative ions are produced from a molecular compound that is dissolved in solution. – Arrhenius acids Return to TOC Copyright © Cengage Learning. All rights reserved 4
  • 5. Section 10.1 Arrhenius Acid-Base Theory Dissociation • The process in which individual positive and negative ions are released from an ionic compound that is dissolved in solution. – Arrhenius Bases Return to TOC Copyright © Cengage Learning. All rights reserved 5
  • 6. Section 10.1 Arrhenius Acid-Base Theory Difference Between Ionization and Dissociation Return to TOC Copyright © Cengage Learning. All rights reserved 6
  • 7. Section 10.2 Brønsted-Lowry Acid-Base Theory • Brønsted-Lowry acid: substance that can donate a proton (H+ ion) to some other substance; proton donor. • Brønsted-Lowry base: substance that can accept a proton (H+ ion) from some other substance; proton acceptor. HCl + H2O → Cl− + H3O+ acid base Return to TOC Copyright © Cengage Learning. All rights reserved 7
  • 8. Section 10.2 Brønsted-Lowry Acid-Base Theory Brønsted-Lowry Reaction Return to TOC Copyright © Cengage Learning. All rights reserved 8
  • 9. Section 10.2 Brønsted-Lowry Acid-Base Theory Acid in Water HA(aq) + H2O(l) H3O+(aq) + A-(aq) acid base conjugate conjugate acid base Return to TOC Copyright © Cengage Learning. All rights reserved 9
  • 10. Section 10.2 Brønsted-Lowry Acid-Base Theory Acid Ionization Equilibrium Return to TOC Copyright © Cengage Learning. All rights reserved 10
  • 11. Section 10.2 Brønsted-Lowry Acid-Base Theory Amphiprotic Substance • A substance that can either lose or accept a proton and thus can function as either a Brønsted-Lowry acid or a Brønsted-Lowry base.  Example: H2O, H3O+ H2O, OH– Return to TOC Copyright © Cengage Learning. All rights reserved 11
  • 12. Section 10.3 Mono-, Di-, and Triprotic Acids Monoprotic Acid • An acid that supplies one proton (H+ ion) per molecule during an acid-base reaction. HA + H2O A− + H3O+ Return to TOC Copyright © Cengage Learning. All rights reserved 12
  • 13. Section 10.3 Mono-, Di-, and Triprotic Acids Diprotic Acid • An acid that supplies two protons (H+ ions) per molecule during an acid-base reaction. H2A + H2O HA− + H3O+ HA− + H2O A2− + H3O+ Return to TOC Copyright © Cengage Learning. All rights reserved 13
  • 14. Section 10.3 Mono-, Di-, and Triprotic Acids Triprotic Acid • An acid that supplies three protons (H+ ions) per molecule during an acid-base reaction. H3A + H2O H2A− + H3O+ H2A− + H2O HA2− + H3O+ HA2− + H2O A3− + H3O+ Return to TOC Copyright © Cengage Learning. All rights reserved 14
  • 15. Section 10.3 Mono-, Di-, and Triprotic Acids Polyprotic Acid • An acid that supplies two or more protons (H+ ions) during an acid-base reaction. • Includes both diprotic and triprotic acids. Return to TOC Copyright © Cengage Learning. All rights reserved 15
  • 16. Section 10.4 Strengths of Acids and Bases Strong Acid • Transfers ~100% of its protons to water in an aqueous solution. • Ionization equilibrium lies far to the right. • Yields a weak conjugate base. Return to TOC Copyright © Cengage Learning. All rights reserved 16
  • 17. Section 10.4 Strengths of Acids and Bases Commonly Encountered Strong Acids Return to TOC Copyright © Cengage Learning. All rights reserved 17
  • 18. Section 10.4 Strengths of Acids and Bases Weak Acid • Transfers only a small % of its protons to water in an aqueous solution. • Ionization equilibrium lies far to the left. • Weaker the acid, stronger its conjugate base. Return to TOC Copyright © Cengage Learning. All rights reserved 18
  • 19. Section 10.4 Strengths of Acids and Bases Differences Between Strong and Weak Acids in Terms of Species Present Return to TOC Copyright © Cengage Learning. All rights reserved 19
  • 20. Section 10.4 Strengths of Acids and Bases Bases • Strong bases: hydroxides of Groups IA and IIA. Return to TOC Copyright © Cengage Learning. All rights reserved 20
  • 21. Section 10.5 Ionization Constants for Acids and Bases Acid Ionization Constant • The equilibrium constant for the reaction of a weak acid with water. HA(aq) + H2O(l) H3O+(aq) + A-(aq) H3O+   A −  Ka =    [ HA ] Return to TOC Copyright © Cengage Learning. All rights reserved 21
  • 22. Section 10.5 Ionization Constants for Acids and Bases Acid Strength, % Ionization, and Ka Magnitude • Acid strength increases as % ionization increases. • Acid strength increases as the magnitude of Ka increases. • % ionization increases as the magnitude of Ka increases. Return to TOC Copyright © Cengage Learning. All rights reserved 22
  • 23. Section 10.5 Ionization Constants for Acids and Bases Base Ionization Constant • The equilibrium constant for the reaction of a weak base with water. B(aq) + H2O(l) BH+(aq) + OH–(aq) BH+  OH−  Kb =    [ B] Return to TOC Copyright © Cengage Learning. All rights reserved 23
  • 24. Section 10.6 Salts • Ionic compounds containing a metal or polyatomic ion as the positive ion and a nonmetal or polyatomic ion (except hydroxide) as the negative ion. • All common soluble salts are completely dissociated into ions in solution. Return to TOC Copyright © Cengage Learning. All rights reserved 24
  • 25. Section 10.7 Acid-Base Neutralization Reactions Neutralization Reaction • The chemical reaction between an acid and a hydroxide base in which a salt and water are the products. HCl + NaOH → NaCl + H2O H2SO4 + 2 KOH → K2SO4 + 2 H2O Return to TOC Copyright © Cengage Learning. All rights reserved 25
  • 26. Section 10.7 Acid-Base Neutralization Reactions Formation of Water Return to TOC Copyright © Cengage Learning. All rights reserved 26
  • 27. Section 10.8 Self-Ionization of Water Self-Ionization • Water molecules in pure water interact with one another to form ions. H2O + H2O H3O+ + OH– • Net effect is the formation of equal amounts of hydronium and hydroxide ions. Return to TOC Copyright © Cengage Learning. All rights reserved 27
  • 28. Section 10.8 Self-Ionization of Water Self-Ionization of Water Return to TOC Copyright © Cengage Learning. All rights reserved 28
  • 29. Section 10.8 Self-Ionization of Water Ion Product Constant for Water • At 24°C: Kw = [H3O+][OH–] = 1.00 × 10–14 • No matter what the solution contains, the product of [H3O+] and [OH–] must always equal 1.00 × 10–14. Return to TOC Copyright © Cengage Learning. All rights reserved 29
  • 30. Section 10.8 Self-Ionization of Water Relationship Between [H3O+] and [OH–] Return to TOC Copyright © Cengage Learning. All rights reserved 30
  • 31. Section 10.8 Self-Ionization of Water Three Possible Situations • [H3O+] = [OH–]; neutral solution • [H3O+] > [OH–]; acidic solution • [H3O+] < [OH–]; basic solution Return to TOC Copyright © Cengage Learning. All rights reserved 31
  • 32. Section 10.8 Self-Ionization of Water Exercise Calculate [H3O+] or [OH–] as required for each of the following solutions at 24°C, and state whether the solution is neutral, acidic, or basic. b) 1.0 × 10–4 M OH– 1.0 × 10–10 M H3O+; basic b) 2.0 M H3O+ 5.0 × 10–15 M OH–; acidic Return to TOC Copyright © Cengage Learning. All rights reserved 32
  • 33. Section 10.9 The pH Concept • pH = –log[H3O+] • A compact way to represent solution acidity. • pH decreases as [H+] increases. • pH range between 0 to 14 in aqueous solutions at 24°C. Return to TOC Copyright © Cengage Learning. All rights reserved 33
  • 34. Section 10.9 The pH Concept Exercise Calculate the pH for each of the following solutions. a) 1.0 × 10–4 M H3O+ pH = 4.00 – 0.040 M OH– pH = 12.60 Return to TOC Copyright © Cengage Learning. All rights reserved 34
  • 35. Section 10.9 The pH Concept Exercise The pH of a solution is 5.85. What is the [H3O+] for this solution? [H3O+] = 1.4 × 10–6 M Return to TOC Copyright © Cengage Learning. All rights reserved 35
  • 36. Section 10.9 The pH Concept pH Range • pH = 7; neutral • pH > 7; basic – Higher the pH, more basic. • pH < 7; acidic – Lower the pH, more acidic. Return to TOC Copyright © Cengage Learning. All rights reserved 36
  • 37. Section 10.9 The pH Concept Relationships Among pH Values, [H3O+], and [OH–] Return to TOC Copyright © Cengage Learning. All rights reserved 37
  • 38. Section 10.10 The pKa Method for Expressing Acid Strength • pKa = –log Ka • pKa is calculated from Ka in exactly the same way that pH is calculated from [H3O+]. Return to TOC Copyright © Cengage Learning. All rights reserved 38
  • 39. Section 10.10 The pKa Method for Expressing Acid Strength Exercise Calculate the pKa for HF given that the Ka for this acid is 6.8 × 10–4. pKa = 3.17 Return to TOC Copyright © Cengage Learning. All rights reserved 39
  • 40. Section 10.11 The pH of Aqueous Salt Solutions Salts • Ionic compounds. • When dissolved in water, break up into its ions (which can behave as acids or bases). • Hydrolysis – the reaction of a salt with water to produce hydronium ion or hydroxide ion or both. Return to TOC Copyright © Cengage Learning. All rights reserved 40
  • 41. Section 10.11 The pH of Aqueous Salt Solutions Types of Salt Hydrolysis • The salt of a strong acid and a strong base does not hydrolyze, so the solution is neutral.  KCl, NaNO3 Return to TOC Copyright © Cengage Learning. All rights reserved 41
  • 42. Section 10.11 The pH of Aqueous Salt Solutions Types of Salt Hydrolysis • The salt of a strong acid and a weak base hydrolyzes to produce an acidic solution.  NH4Cl NH4+ + H2O → NH3 + H3O+ Return to TOC Copyright © Cengage Learning. All rights reserved 42
  • 43. Section 10.11 The pH of Aqueous Salt Solutions Types of Salt Hydrolysis • The salt of a weak acid and a strong base hydrolyzes to produce a basic solution.  NaF, KC2H3O2 F– + H2O → HF + OH– C2H3O2– + H2O → HC2H3O2 + OH– Return to TOC Copyright © Cengage Learning. All rights reserved 43
  • 44. Section 10.11 The pH of Aqueous Salt Solutions Types of Salt Hydrolysis • The salt of a weak acid and a weak base hydrolyzes to produce a slightly acidic, neutral, or slightly basic solution, depending on the relative weaknesses of the acid and base. Return to TOC Copyright © Cengage Learning. All rights reserved 44
  • 45. Section 10.11 The pH of Aqueous Salt Solutions Neutralization “Parentage” of Salts Return to TOC Copyright © Cengage Learning. All rights reserved 45
  • 46. Section 10.12 Buffers Key Points about Buffers • Buffer – an aqueous solution containing substances that prevent major changes in solution pH when small amounts of acid or base are added to it. • They are weak acids or bases containing a common ion. • Typically, a buffer system is composed of a weak acid and its conjugate base. Return to TOC Copyright © Cengage Learning. All rights reserved 46
  • 47. Section 10.12 Buffers Buffers Contain Two Active Chemical Species 1. A substance to react with and remove added base. 2. A substance to react with and remove added acid. Return to TOC Copyright © Cengage Learning. All rights reserved 47
  • 48. Section 10.12 Buffers Adding an Acid to a Buffer Return to TOC Copyright © Cengage Learning. All rights reserved 48
  • 49. Section 10.12 Buffers Buffers Return to TOC Copyright © Cengage Learning. All rights reserved 49
  • 50. Section 10.12 Buffers Addition of Base [OH– ion] to the Buffer HA + H2O H3O+ + A– • The added OH– ion reacts with H3O+ ion, producing water (neutralization). • The neutralization reaction produces the stress of not enough H3O+ ion because H3O+ ion was consumed in the neutralization. • The equilibrium shifts to the right to produce more H3O+ ion, which maintains the pH close to its original level. Return to TOC Copyright © Cengage Learning. All rights reserved 50
  • 51. Section 10.12 Buffers Addition of Acid [H3O+ ion] to the Buffer HA + H2O H3O+ + A– • The added H3O+ ion increases the overall amount of H3O+ ion present. • The stress on the system is too much H3O+ ion. • The equilibrium shifts to the left consuming most of the excess H3O+ ion and resulting in a pH close to the original level. Return to TOC Copyright © Cengage Learning. All rights reserved 51
  • 52. Section 10.13 The Henderson-Hasselbalch Equation Henderson-Hasselbalch Equation − A  pH = pK a + log   [ HA ] Return to TOC Copyright © Cengage Learning. All rights reserved 52
  • 53. Section 10.13 The Henderson-Hasselbalch Equation Exercise What is the pH of a buffer solution that is 0.45 M acetic acid (HC2H3O2) and 0.85 M sodium acetate (NaC2H3O2)? The Ka for acetic acid is 1.8 × 10–5. pH = 5.02 Return to TOC Copyright © Cengage Learning. All rights reserved 53
  • 54. Section 10.14 Electrolytes • Acids, bases, and soluble salts all produce ions in solution, thus they all produce solutions that conduct electricity. • Electrolyte – substance whose aqueous solution conducts electricity. Return to TOC Copyright © Cengage Learning. All rights reserved 54
  • 55. Section 10.14 Electrolytes Nonelectrolyte – does not conduct electricity • Example: table sugar (sucrose), glucose Return to TOC Copyright © Cengage Learning. All rights reserved 55
  • 56. Section 10.14 Electrolytes Strong Electrolyte – completely ionizes/dissociates • Example: strong acids, bases, and soluble salts Return to TOC Copyright © Cengage Learning. All rights reserved 56
  • 57. Section 10.14 Electrolytes Weak Electrolyte – incompletely ionizes/dissociates • Example: weak acids and bases Return to TOC Copyright © Cengage Learning. All rights reserved 57
  • 58. Section 10.15 Equivalents and Milliequivalents of Electrolytes Equivalent (Eq) of an Ion • The molar amount of that ion needed to supply one mole of positive or negative charge. 1 mole K+ = 1 equivalent 1 mole Mg2+ = 2 equivalents 1 mole PO43– = 3 equivalents Return to TOC Copyright © Cengage Learning. All rights reserved 58
  • 59. Section 10.15 Equivalents and Milliequivalents of Electrolytes Milliequivalent 1 milliequivalent = 10–3 equivalent Return to TOC Copyright © Cengage Learning. All rights reserved 59
  • 60. Section 10.15 Equivalents and Milliequivalents of Electrolytes Concentrations of Major Electrolytes in Blood Plasma Return to TOC Copyright © Cengage Learning. All rights reserved 60
  • 61. Section 10.15 Equivalents and Milliequivalents of Electrolytes Exercise The concentration of Ca2+ ion present in a sample is 5.3 mEq/L. How many milligrams of Ca2+ ion are present in 180.0 mL of the sample? 19 mg Ca2+ ion ( )( )( )( )( )( ) 2+ 2+ ( 180 mL ) 1L 1000 mL 5.3 mEq 1L 1Eq 1000 mEq 1 mol Ca 2 Eq Ca 2+ 40.08 g Ca 1 mol Ca 2+ 1000 mg 1g = 19 mg Ca 2+ ion Return to TOC Copyright © Cengage Learning. All rights reserved 61
  • 62. Section 10.16 Acid-Base Titrations • A neutralization reaction in which a measured volume of an acid or a base of known concentration is completely reacted with a measured volume of a base or an acid of unknown concentration. • For a strong acid and base reaction: H+(aq) + OH–(aq) → H2O(l) Return to TOC Copyright © Cengage Learning. All rights reserved 62
  • 63. Section 10.16 Acid-Base Titrations Titration Setup Return to TOC Copyright © Cengage Learning. All rights reserved 63
  • 64. Section 10.16 Acid-Base Titrations Acid-Base Indicator • A compound that exhibits different colors depending on the pH of its solution. • An indicator is selected that changes color at a pH that corresponds as nearly as possible to the pH of the solution when the titration is complete. Return to TOC Copyright © Cengage Learning. All rights reserved 64
  • 65. Section 10.16 Acid-Base Titrations Indicator – yellow in acidic solution; red in basic solution Return to TOC Copyright © Cengage Learning. All rights reserved 65
  • 66. Section 10.16 Acid-Base Titrations Concept Check For the titration of sulfuric acid (H2SO4) with sodium hydroxide (NaOH), how many moles of sodium hydroxide would be required to react with 1.00 L of 0.500 M sulfuric acid to reach the endpoint? 1.00 mol NaOH Return to TOC Copyright © Cengage Learning. All rights reserved 66

Editor's Notes

  1. K w = [H 3 O + ][OH – ] = 1.00 × 10 – 14 1.00 × 10 – 14 = [H 3 O + ]( 1.0 × 10 –4 M) = 1.0 × 10 –10 M H 3 O + ; basic 1.00 × 10 – 14 = (2.0)[ OH – ] = 5.0 × 10 –15 M OH – ; acidic
  2. pH = –log[H 3 O + ] a) pH = –log[H 3 O + ] = –log( 1.0 × 10 –4 M ) = 4.00 b) K w = [H 3 O + ][OH – ] = 1.00 × 10 – 14 = [H 3 O + ](0.040 M) = 2.5 × 10 – 13 M H 3 O + pH = –log[H 3 O + ] = –log( 2.5 × 10 – 13 M ) = 12.60
  3. [H 3 O + ] = 10^–5.85 = 1.4 × 10 –6 M
  4. p K a = –log K a = –log(6.8 × 10 –4 M ) = 3.17
  5. See notes on slide 1.
  6. See notes on slide 1.
  7. See notes on slide 1.
  8. See notes on slide 1.
  9. See notes on slide 1.
  10. See notes on slide 1.
  11. pH = –log K a + log([ C 2 H 3 O 2 – ] / [H C 2 H 3 O 2 ]) = –log( 1.8 × 10 –5 ) + log(0.85 M / 0.45 M) = 5.02
  12. (180.0 mL)(1 L/1000 mL)(5.3 mEq/L)(1 Eq/1000 mEq)(1 mol Ca 2+ /2 Eq Ca 2+ )(40.08 g/mol)(1000 mg/1g) = 19 mg Ca 2+ ion
  13. 1.00 mol of sodium hydroxide would be required.