16.S: Acid–Base Equilibria (Summary)
- Page ID
- 70556
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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