Scruggs Chap 16b

Scruggs Chap 16b - 16.6 Weak Acids q Except for the strong...

Info iconThis preview shows page 1. Sign up to view the full content.

View Full Document Right Arrow Icon
This is the end of the preview. Sign up to access the rest of the document.

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 ...
View Full Document

This note was uploaded on 04/06/2011 for the course CHM 116 taught by Professor Unknown during the Spring '08 term at ASU.

Ask a homework question - tutors are online