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AP Chemistry

AP Chemistry. Acids and Bases. Aqueous Equilibria: Acids and Bases. Arrhenius Acids and Bases Acids cause [H+] to increase, bases cause [OH-] to increase Bronsted-Lowry Acids and Bases H + /proton Donor (acid) and H +/ proton Acceptor (base) Lewis Acid and Bases Acids accept electron pair

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AP Chemistry

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  1. AP Chemistry Acids and Bases

  2. Aqueous Equilibria: Acids and Bases Arrhenius Acids and Bases Acids cause [H+] to increase, bases cause [OH-] to increase Bronsted-Lowry Acids and Bases H+ /proton Donor (acid) and H+/proton Acceptor (base) Lewis Acid and Bases Acids accept electron pair Bases donate electron pair

  3. Acid and Base Strengths • Based on extent of dissociation. • Strong Acids Dissociate nearly 100% • If HA  H+ + A- • A- is a very weak base. (the conjugate base) • Acid and Conjugate Base explanation of strength. Pair of substances differing only by H+ • HF(aq) + H2O(l) <==> H3O+(aq) + F-(aq) • acid 1 base 2 acid 2 base 1 • H3O+(aq) + OH-(aq) <==> H2O(l) + H2O(l) • acid 1 base 2 acid 2 base 1

  4. Acid and Base Strengths • Taken from State University of West Georgia Chemistry Dept.

  5. Acid and Base Strengths • Taken from State University of West Georgia Chemistry Dept.

  6. Hydronium Ions

  7. Hydronium Ions H5O2+

  8. Dissociation of Water • H2O + H2O <---> H3O1+ + OH1- • The equilibrium expression is products over reactants. • K = [H3O1+] [OH1-] / [H2O] [H2O] • The molarity for the water is a constant at any specific temperature. So • K [H2O] [H2O] = [H3O1+] [OH1-] • The quantity on the right hand side of the equation is formally defined as Kw. The numerical vale for Kw is different at different temperatures. • At 25oC Kw = 1.014 x 10-14 • Kw = K[H2O] [H2O] or Kw = [H3O1+] [OH1-]

  9. Dissociation of Water • Equilibrium constants exist then for both acid dissociation and base. (Ka and Kb) • The higher the Ka, the stronger the acid and the higher the Kb, the stronger the base. • Ka and Kb are related by the previous equation. • Kw = KaKb

  10. Dissociation of Water • As Ka gets larger the strength of the acid gets higher, but Kb must fall.  Therefore the stronger the acid, the weaker the conjugate base. • It can now be said that the conjugate base (acid) of a weak acid (base) is a weak base (acid) and the conjugate base (acid) of a strong acid (base) is a worthless base (acid). • The strength of an acid/base is usually given as a pKa value. As pKa is inversely related to Ka, the higher the Ka (the stronger the acid), the lower the pKa value. The same is true of bases.

  11. Calculating pH -log [H+] Power of Hydronium (Hydrogen) P[OH-] = - log [OH]

  12. The pH Scale

  13. pH in Solutions of Strong Acids and Strong Bases • Strong acids • Certain acids are known as strong acids. These are acids that fully ionize when placed in water: • HA + H2O  A- + H3O+ • Goes to completion and thus • Ka = [A-][H3O+]/[HA] = infinity • Some common strong acids are: • HCl, hydrochloric acid • HBr, hyrdobromic acid • HI, hydroiodic acid • H2SO4, sulfuric acid • HNO3, nitric acid • HClO4, perchloric acid

  14. pH in Solutions of Strong Acids and Strong Bases • Strong Bases • Certain bases are known as strong bases. These are bases that fully ionize when placed in water. • Some common strong bases are: • LiOH, lithium hydroxide • NaOH, sodium hydroxide • KOH, potassium hydroxide • Ca(OH)2, calcium hydroxide • Sr(OH)2, strontium hydroxide • Ba(OH)2, barium hydroxide • Alkaline earth oxides. • Lime (CaO)

  15. Equilibrium in Solutions of Weak Acids • HA(aq) + H2O(l)  A-(aq) + H3O+(aq) • The equilibrium constant for a weak acid is • Ka = [H3O+][A-]/[HA] • For a weak acid then Ka << 1 • For a strong acid Ka >> 1 • A common way to express the strength of an acid is the pKa, which is similar in form to the pH • pKa = -log10Ka

  16. Calculating Equilibrium Concentrations in Solutions of Weak Acids • Principle Reaction vs Subsidiary Reactions. • If one of the equilibrium reactions is less than 100 x the extent of the other. • Always check • H2O(l) + H2O(l) <---> H3O+(aq) + OH-(aq) • Kw = 1.0 x 10-14

  17. Percent Dissociation in Solution of Weak Acids • Percent dissociation = [HA] dissociated / [HA] initial x 100%

  18. More Discussion • Acid • HA + S  HS+ + A- • Acid Solvent Conjugate acid Conjugate base • Base • HB+ + S  HS+ + B • Conjugate acid Solvent Acid Base • B + H2O  HB+ + OH- • Kb =

  19. More Discussion • Kw = Ka x Kb • Or pKa + pKb = pKw • Carbonic acid (H2CO3) (Data in H2O) • 1. Ka = 4.3 x 10-7 pKa = 6.37 • 2. Ka = 5.61 x 10 -11 pKa = 10.25 • Explain what happens when the Ka of an acid is smaller than the Ka for H2O.

  20. Polyprotic Acids • A polyprotic acid is one that has multiple ionizable protons, such as H2SO4 or H3PO4. • Each proton has its own equilibrium constant Ka. For example, for a diprotic acid H2A, • H2A(aq)  H+(aq) + HA-(aq)   Ka1 = [H+][HA-]/[H2A] • HA-(aq)  H+(aq) + A-2(aq)     Ka2 = [H+][A-2]/[HA-] • In general, Ka1 >> Ka2 >> Ka3. • You can compute the K for the total ionization of the acid. If you add the above equations. • H2A (aq)  2H+(aq) + A-(aq)          Ktotal = Ka1*Ka2

  21. Polyprotic AcidsTaken from University of Alberta chemistry dept. • Ionization Constants of Aqueous Polyprotic Acids • Common Formula Dissociation Constant pKa • arsenic acid H3AsO4 K1 = 5.65 x 10-3 2.248 - H2AsO4- K2 = 1.75 x 10-7 6.757 - HAsO42- K3 = 2.54 x 10-12 11.596 • boric acid H3BO3 K1 = 5.78 x 10-10 9.238 • carbonic acid H2CO3 K1 = 4.35 x 10-7 6.361 - HCO3- K2 = 4.69 x 10-11 10.329 • chromic acid H2CrO4 K1 = 3.55 -0.550 • - HCrO4- K2 = 3.36 x 10-7 6.473 • citric acid HOC(CH2COOH)3 K1 = 7.42 x 10-4 3.130 - - K2 = 1.75 x 10-5 4.757 - - K3 = 3.99 x 10-6 5.602 • EDTA C2H4N2(CH2COOH)4 K1 = 9.81 x 10-3 2.008 - - K2 = 2.08 x 10-3 2.683 - - K3 = 7.98 x 10-7 6.098 - - K4 = 6.60 x 10-11 10.181

  22. Polyprotic AcidsTaken from University of Alberta chemistry dept. • Common Formula Dissociation Constant pKa • glycinium ion H3NCH2COOH+ K1 = 4.47 x 10-3 2.350 -(glycine) H2NCH2COOH K2 = 1.67 x 10-10 9.778 • hydrogen sulfide H2S K1 = 1.02 x 10-7 6.992 - HS- K2 = 1.22 x 10-13 12.915 • oxalic acid HOOCCOOH K1 = 5.40 x 10-2 1.268 - HOOCCOO- K2 = 5.23 x 10-5 4.282 • phthalic acid C6H4(COOH)2 K1 = 1.13 x 10-3 2.946 - - K2 = 3.90 x 10-6 5.409 • phosphoric acid H3PO4 K1 = 7.11 x 10-3 2.148 - H2PO4- K2 = 6.23 x 10-8 7.206 - HPO42- K3 = 4.55 x 10-13 12.342 • succinic acid C(CH2)2COOH K1 = 6.21 x 10-5 4.207 - HOOC(CH2)2COO- K2 = 2.31 x 10-6 5.636 • sulfuric acid H2SO4 K1 > 1 negative - HSO4- K2 = 1.01 x 10-2 1.994 • sulfurous acid H2SO3 K1 = 1.71 x 10-2 1.766 - HSO3- K2 = 5.98 x 10-8 7.223

  23. 15.11Polyprotic Acids

  24. Equilibria in Solutions of Weak Bases • Remember Kw = KaKb

  25. Relation Between Ka and Kb • HA(aq) + H2O(l)  H3O+(aq) + A-(aq) • Ka • A-(aq)+ H2O(l)  HA(aq) + OH-(aq) • Kb • Kw = KaKb = 1.0 x 10-14 • Ka = Kw / Kb • Kb = Kw / Ka • Knet = K1 x K2 x K3……

  26. Acid/Base

  27. Acid/Base

  28. Acid-Base Properties of Salts Acid Strength Strong Weak Strong Base Strength Weak Resulting Salt Solution

  29. Acid-Base Properties of Salts

  30. Acid-Base Properties of Salts • Example #1 • NaOH(aq) + HCl(aq)  NaCl(aq) + H2O Strong Base Strong Acid Neutral Salt

  31. Acid-Base Properties of Salts • Example #2 • NaOH(aq) + HF(aq)  NaF(aq) + H2O Strong Base Weak Acid Basic Salt

  32. Acid-Base Properties of Salts • Example #3 • NH3(aq) + HCl(aq)  NH4Cl(aq) Weak Base Strong Acid Acidic Salt

  33. Acid-Base Properties of Salts • Example #4 • NH3(aq) + CH3COOH(aq)  NH4OOCCH3(aq) Weak Base Weak Acid ? Salt • Compare Ka to Kb • Ka = 5.6 x 10-10 of NH4+ • Kb = 5.7 x 10-10 of -OOCCH3 • Salt is Neutral or ?

  34. Acid-Base Properties of Salts • Example #5 • 2NH3(aq) + H2CO3(aq)  (NH4)2CO3(aq) Weak Base Weak Acid Acidic Salt • Compare Ka to Kb • Ka = 5.6 x 10-10 of NH4+ • Kb = 1.8 x 10-4 of CO32- • Salt is Basic

  35. Factors That Affect Acid Strength • HA  H+ + A- • Extent of dissociation depends on H-A bond strength and Electronegativity (or stability of negative charge) on A. • This explanation works for Halogen acids, Organic oxoacids, or Inorganic oxoacids. • Oxidation # of the Halide is not necessary.

  36. Lewis Acids and Bases • Bronsted-Lowry Acid • Proton Donor • Lewis Acid • Electron Pair Acceptor • Bronsted-Lowry Base • Proton Acceptor • Lewis Base • Electron Pair Acceptor

  37. Lewis Acids and Bases • Lewis Acid • Anything with a vacant valence orbital • Charged or Neutral • Other Examples • Fe+3(aq) + 6CN-(aq) -> Fe(CN)6-3(aq) • Cu+2(aq) + 4NH3(aq) -> Cu(NH3)4+2(aq)

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