16.S: Acid–Base Equilibria (Summary)

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
  • k_w = 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 K_w = 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)
  • 
  • K_a = acid - dissociation constant
  • The lager the K_a the stronger the acid
  • K_a 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 K_a1, K_a2, etc…
  • K_a values usually differ by 10³

16.7: Weak Bases

• • base-dissociation constant, K_b
  • 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 K_a and K_b

• • 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
  • K_a x K_b = [H⁺][OH^-] = K_w
  • Acid-dissociation constant times base-dissociation constant equals the ion-product constant for water
  • K_a x K_b = K_w = 1.0 x 10^-14
  • pK_a x pK_b = pK_w = 14; (pK_a= -log K_a and pK_b = -log K_b)

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 K_a and K_b
  • 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 K_a increases with increasing charge and decreasing radius of ion