Unformatted text preview: 16.6 Weak Acids
q Except for the strong acids, most acids do not ionize completely. These acids are called weak acids. q HF(aq) + H2O(l) 6 H3O+(aq) + F-(aq) Weak Acids
q HF(aq) + H2O(l) 6 H3O+(aq) + F-(aq) q What is the equilibrium constant expression? q We call K for an acid the acid ionization constant, K . H2O(l) is omitted. [H3O+][F-] K = ---------- a Some Weak Acids Weak Acids
q Note that polyprotic acids have more than one Ka value and each successive value decreases, often by a factor of about 10-4 or 10-5.
q H2CO3 + H2O 6 H3O+ + HCO3 q HCO3- + H2O 6 H3O+ + CO326 Ka1 = 4.45 x 10-7 Ka2 = 4.7 x 10-11 q Which ionizes more completely:
H2CO3 or HCO3-? Some Polyprotic Weak Acids Calculation of pH for Weak Acids
q Example: What is the pH of a 0.100 M HF solution? Ka = 7.0 x 10-4 q HF(aq) + H2O(l) 6 H3O+(aq) + F-(aq) q Ka = [H3O+][F-]/[HF] = 7.0 x 10-4
q Substance Initial Change Equilibrium HF 0.100 -x 0.100 - x H3O+ 0 +x x F0 +x x pH of 0.100 M HF
q (x)(x)/(0.100 - x) = 7.0 x 10-4 q Assume x << 0.100 q x2/0.100 = 7.0 x 10-4 q x2 = 7.0 x 10-5 q x = 8.4 x 10-3, which is 8% of initial concentration, so assumption was not very good, and we must solve the quadratic equation. q (x)(x)/(0.100 - x) = 7.0 x 10-4 q x2 = 7.0 x 10-5 - 7.0 x 10-4 x pH of 0.100 M HF
q x2 + 7.0 x 10-4 x - 7.0 x 10-5 = 0
q -7.0 x 10-4 + [(7.0 x 10-4)2 - (4)(1)(-7.0 x 10-5)]1/2 x= ------------------------------------ -- (2)(1) q x = 0.00802 [H3O+] = [F ] = 8.02 x 10 M q
-3 What is the pH of a 1.00M HF solution?
25% 25% 25% 25% Ka = 7.0 x 10-4 1. 2. 3. 4. 1.00 1.58 4.27 7.94
1. 2. 3. 4. Percent Ionization
q What is the percent ionization of HF in a 1.0 M HF solution? 1 pH and Ka
q Should be able to work back from a pH value to a value of Ka. Generally simpler than calculating pH. q The pH of 0.500 M HNO2 is 1.827. What is Ka of HNO2? q [H3O+] = 10-1.827 = 0.0149 M q [NO2-] = [H3O+] = 0.0149 M q [HNO2] = 0.500 - 0.0149 = 0.485 M q Ka = [H3O+][NO2-]/[HNO2] q Ka = (0.0149)2/0.485 = 4.58 x 10-4 1 Which of these lists all of the substances present in a solution of H2CO3, in order of concentration? (high to low) 1. HCO3-, CO3217% 17% 17% 17% 17% 2. CO32-, HCO33. H2CO3, HCO3-, CO324. H2CO3, HCO35. HCO3-, H2CO3 6. CO32- ,HCO3-, H2CO3,
1 2 3 4 5 17% 6 1 16.7 Weak Bases
q Concepts and calculations are similar to those for weak acids. q General Eq: q B(aq) + H2O(l) 6 BH+(aq) + OH-(aq) q Kb = [BH+][OH-]/[B] q Most common weak bases are amines (NH3) or conjugate bases of weak acids. q See Table 16.4 for some representative values of K . 1 Some Weak Bases 1 Weak Bases
q Calculations are the same as for weak acids, except we solve for [OH-] and pOH, then convert to pH by using pKw. q What is the pH of a 0.100 M NH3 solution? q (Kb = 1.76 x 10-5 ) q For 0.100 M NH3, pOH = 2.877 and pH = 11.123 q How does this compare to 0.10 M NaOH? 1 16.8 Relationship Between Ka and Kb
q Ka and Kb for a conjugate acid-base pair are related by Kw Ka Kb = Kw 6 q NH3(aq) + H2O(l) 6 NH4+(aq) + OH-(aq) K = [NH ][OH ]/[NH ] q
b 4+ - 6
3 1 Weak Acid and Base Constants
q Kb = [NH4+][OH-]/[NH3] q Ka = [NH3][H3O+]/[NH4+] q KaKb = = [OH-][H3O+] = Kw
a b w q 1 Weak Acid and Base Constants
q If we know the K for an acid or base we can calculate the K for its conjugate: q Kb(NH3) = 1.76 x 10-5 q What is Ka for NH4+? q Ka = Kw/Kb = (1 10-14) / (1.76 x 10-5) q Ka = 5.68 10-10 1 Weak Acid and Base Constants
q Group Work q The Ka for HF is 6.8x10-4 q The Ka for HCN is 4.9x10-10 q Rank the following anions in order of increasing base strength: F-, CN-, or Cl- 1 16.9 Acid-Base Properties of Salt Solutions
q Salts dissolve in water to form ions q MX(s) M+(aq) + X-(aq) q In the case of NaCl, Na+ and Cl- are so weak that they will not react with water. This is because they are conjugates of a strong base and a strong acid. q NaOH Na+(aq) + OH-(aq) q HCl + H2O H3O+(aq) + Cl-(aq) 2 16.9 Acid-Base Properties of Salt Solutions
q Salts dissolve is water to form ions q MX(s) M+(aq) + X-(aq) q NaF dissolves in water to form a basic solution because F- is the conjugate of the weak acid HF: (HF + H2O 6 H3O+(aq) + F-(aq)) F- reacts with water as a weak base: F- + H2O(aq) 6 HF(aq) + OH6 2 16.9 Acid-Base Properties of Salt Solutions
q Hydrolysis: - The reaction of a cation or anion with water to form H3O+ or OH- ions. q In which of the following solutions should hydrolysis occur? KBr(aq) LiCN(aq) NH4CN(aq) NaNO3(aq) 2 16.9 Acid-Base Properties of Salt Solutions
q Cation/Anion from: q Strong base, q Strong base, strong acid no hydrolysis pH = 7 weak acid anion hydrolysis pH > 7 NaCl LiCN NH4Cl NH4CN q Weak base, q Weak base, strong acid cation hydrolysis pH < 7 weak acid cation and anion hydrolysis 2 Strongest Acid HClO Base ClO Weakest acids H SO 4 4 HSO- bases HI 2 4 I 4 - HBr Br - HCl Cl - 2 Group Work
q Which of the following salts will change the pH of a solution? NaCl KNO2 LiCN KClO4 Na2CO3 NH4Cl CaBr2 2 Of the following salts, which will change the pH of a solution? 20% 20% 20% 20%
NaCl, KNO2 ,KClO4 Na2CO3,NH4Cl 20% 1. KNO2, KClO4 2. NaCl, KClO4 3. KNO2, Na2CO3 4. Na2CO3, NaCl 5. Na2CO3, KNO2, NH4Cl
1 2 3 4 5 2 16.10 Acid-Base Behavior and Chemical Structure
Factors That Affect Acid Strength
Consider the molecule H-X. For this substance to be an acid: q H-X bond must be polar with H + and X (so X can't be a metal). q the H-X bond must be weak enough to be broken. q the conjugate base, X, must be stable. 2 Structure and Bonding Effects on Acid Strength
q Binary acids: Ka increases with electronegativity and size of X. q These factors don't always work in the same direction. q Electronegativity is most important when comparing X in same period (Ka: HCl>H2S). q Size becomes most important when comparing X in same group (Ka: HCl>HF). 2 Binary Acids 2 Structure and Bonding Effects on Acid Strength
q Oxoacids, EOnH q For constant E, an increase in oxidation number results in an increase in Ka. q Greater charge of E in E-O-H pulls electrons from O-H and ionizes this bond more. q +1 +3 +5 +7 HOCl < HOClO < HOClO2 < HOClO3 3 Structure and Bonding Effects on Acid Strength
q When E is not constant, Ka varies with electronegativity of E.
q H3BO3 H2CO3 HNO3 large H3PO4 8 x 10-3 electronegativity increases H3AsO4 qElectronegativity is more important-3than 6 x 10 size of E. (in oxyacids) 6 x 10-10 4 x 10-7 3 Effect of Electronegativity
Compare HOCl to HOI 3 Structure and Bonding Effects on Acid Strength
q Polyprotic acids: successive Ka values decrease by about 105. H3PO4 7.5 x 10-3 H2PO4HPO426.2 x 10-8 2.2 x 10-13 q It is easier to pull H+ away from a neutral molecule than from an anion and the difficulty increases with increasing negative charge. 3 Group Work
q In each pair, which is the strongest acid? q H2S or H2O q HNO3 or HNO2 q HBrO3 or HClO3 3 Rank these oxyacids from weakest to strongest.
H2SO3, HClO3, HClO2
25% 25% 25% 25% 1. HClO3, H2SO3, HClO2 2. HClO2, H2SO3, HClO3 3. H2SO3, HClO3, HClO2 4. HClO3, HClO2 H2SO3,
1 2 3 4 3 Oxides and Hydroxides
q Oxides and hydroxides also act as acids or bases. They are the anhydrides (without water) of acids or bases. H2SO4 - H2O = SO q
3 2H3PO4 - 3H2O = P O (or P4O10) 3 Oxides and Hydroxides
q Metal oxides are usually bases. metal oxide + water hydroxide base q Nonmetal oxides are usually acids. nonmetal oxide + water oxoacid 3 Why are metal oxides bases?
q In both metal and nonmetal oxides, hydration produces E-O-H, which can either be an oxoacid or a hydroxide base. q The difference arises from the identity of the bond that is most ionic (greatest electroneg. difference, ), E-O or O-H. q Base: E-O-H(aq) E+(aq) + OH-(aq) q Acid: E-O-H(aq) EO-(aq) + H+(aq) 3 Electronegativity Differences
q Metals have low ; nonmetals have high 3 Periodicity of Acidity
q Acid character of oxides increases from left to right and bottom to top of the periodic table.
acidic basic 4 16.11 Lewis Acids and Bases
q Proton transfer also involves transfer of an electron pair H3N: + H2O:H+ H3N:H+ + H2O q Lewis base: electron pair donor q Lewis acid: electron pair acceptor q Useful for systems that don't have protons q acid + base: acid:base (adduct or coordination compound or complex ion) BF3 + :NH3 6 F3B:NH3 4 Lewis Acids and Bases
q Identify the Lewis acid and the Lewis base: Co2+ + 4 ClCoCl42Cu2+ + 4 NH3 6 Cu(NH3)42+ q The metal ion in each case is the acid and the nonmetal species is the base. The products are called complex ions. 4 Hydrolysis of Metal Ions (Lewis Acid-Base Interaction)
q Ions are often modified when dissolved in water. q Fe(NO3)3.6H2O contains pink Fe(H2O)63+ (due to hydration)
q Solutions may hydrolyze to give yellow Fe(H2O)5OH2+ or even reddish brown Fe(H2O)3(OH)3 4 Ions in Water 4 Hydrolysis of Metal Ions
q Hydrolysis of metal ions is more important for more highly charged ions - the stronger interactions allow for deprotonation of the hydrated H2O. 4 Hydrolysis of Metal Ions
q Highly charged metal ions (> +3) cause pH shifts due to hydrolysis: Fe(H2O)63+ + H2O Fe(H2O)5OH2+ + H3O+ 4 Hydrolysis of Metal Ions
q Hydrolysis of metal ions is observed for: Cations with charge > +3 Transition metal +2 ions Some post-transition metal ions with high charge q Common for: q Fe3+, Cr3+, Al3+, Zn2+, Cu2+, Bi3+, Pb4+ 4 Hydrolysis of Metal Ions
q See table below for values of Ka for metal ions.
Na+ 95 pm 3.3 x 10-15 Li+ Be2+ Mg2+ Ba2+ Cr 60 pm 31 pm 65 pm 135 pm 69 pm 1.5 x 10-14 3.2 x 10-7 3.8 x 10-12 1.5 x 10-14 9.8 x 10 4 Hydrolysis of Metal Ions
q Which of the following will dissolve in water to form acidic solutions? q CaCl2 q FeCl3 q AgNO3 q Pb(NO3)4 4 ...
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