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Acids and Bases

Acids and Bases. Chapter 17. Why are lemons sour?. What about other acids?. What about bases?. Definitions . Br ønsted-Lowry: An acid is anything that can donate a proton HCl (aq) + H 2 O (l)  Cl - (aq) + H 3 O + (aq) A base is anything that can accept a proton

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Acids and Bases

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  1. Acids and Bases Chapter 17

  2. Why are lemons sour?

  3. What about other acids?

  4. What about bases?

  5. Definitions • Brønsted-Lowry: • An acid is anything that can donate a proton HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq) • A base is anything that can accept a proton • Water (above) acts as a base • H3O+ = hydronium ion • HCl = acid, Cl- = conjugate base • Conjugate acid-base pair • H2O = base, H3O+ = conjugate acid • Conjugate acid-base pair

  6. A problem • Write the conjugate bases for the acids and the conjugate acids for the bases • Acids Bases • HIO3 S2- • HCO3- CN- • CH3COOH NH3

  7. Some more traits • Species donates 1 proton = monoprotic acid • 2 protons = diprotic acid • Accepts 1 proton = monoprotic base • Accepts 2 protons = diprotic base • Give a triprotic acid

  8. Lewis acids and bases • The same dude who gave us Lewis structures • Lewis acid = lone pair acceptor • Lewis base = lone pair donor • Form coordinate covalent bonds or adducts • NH3 + BH3

  9. Coordination complexes • = adducts • Water, ammonia, ions, molecules, etc. (ligands) give lone pairs to (usually transition) metal cations in soln • Ex: Cu2+(aq) + 4NH3(aq) [Cu(NH3)4]2+(aq) • Very colorful solns • Chem 133

  10. Water’s uniqueness • Can act as both a base and an acid • Amphiprotic H2O(l) + H2O(l) H3O+(aq)+ OH-(aq) • Water autoionizes

  11. pH • = power of hydrogen • Measurement of protons/hydronium ions in soln • Based on -logarithm (log10) •  pH = -log[H3O+] • Likewise • pOH = -log[OH-] • (10-pH = [H3O+] & 10-pOH = [OH-])

  12. pH scale • So pH + pOH = 14.00 = pKw • Where Kw = [H3O+][OH-] = 1.0 x 10-14 • 7.00 = neutrality • <7.00  acidity (increase going to 0) • >7.00  basicity (increases going to 14) • At neutrality, [H3O+] = [OH-] = 1.0 x 10-7 • Is [OH-] more concentrated under acidic or basic conditions?

  13. Before you work on next slide! • Sig figs in logs: • Result has the same number of decimal places in the input has sig figs: Log (1.00  10–5) = -5.000 3 sig figs 3 dec place • Sig figs in antilogs: • Result has the same number of sig figs as the number of decimal places in the input: 10–6.00 = 1.0  10–6 2 dec place 2 sig figs

  14. pH calculation problems

  15. Equilibria for acids: HA(aq) + H2O(l) H3O+(aq) + A-(aq) • As Ka increases what happens to the acid dissociation? Does it increase or decrease? • It increases • A stronger acid • If Ka < 1.0, weak acid

  16. Strong acids • HCl • HBr • HI • H2SO4 • HNO3 • HClO4 • HClO3

  17. Equilibria for bases:B(aq) + H2O(l) BH+(aq) + OH-(aq) • As Kb increases so does base strength • If Kb < 1.0 • Weak base

  18. Strong bases • Metal (I) hydroxides • Ba(OH)2

  19. Comparing acid/base strengths • Can use a logarithmic scale pKa = -log Ka • As Ka increases acid strength increases • But what about pKa? Does it increase or decrease as Ka increases? • What can we say about acid strength and pKa?

  20. Conversely pKb = -log Kb • Likewise, as pKb decreases base strength increases

  21. The relation between Ka and Kb HCN(aq) + H2O(l) H3O+(aq) + CN-(aq) Ka = 4.0 x 10-10 CN-(aq) + H2O(l)  HCN(aq) + OH-(aq) Kb = 2.5 x 10-5 Net rxn: 2H2O(l)  H3O+(aq) + OH-(aq) Knet = 1.0 x 10-14 = Kw

  22. Hence • And pKw = pKa + pKb = 14.00

  23. Problem • What is the Ka and pKa of a 0.10 M soln of chloroacetic acid with a pH of 1.95?

  24. Solution

  25. Problem • What is the pH of a 0.237 M solution of benzoic acid? The pKa is 4.19.

  26. Solution

  27. Problem • What is the pH of a 1.00 M solution of Sodium Acetate? (pKb = 5.6 x 10-10) Why pKb and not pKa?

  28. Solution

  29. Another problem • What is the pH of a solution that is made from 250.0 mL of 0.250 M KOH and 150.0 mL of 0.0125 M HBr?

  30. Solution

  31. Which side will be favored? • Rxns proceed from stronger acid/base to weaker acid/base • Predict the products and the direction of arrow for the following reaction: • NH3(aq) + HCO3- • Use table 17.3 on pg. 808

  32. Solution

  33. More problems • For the following write the complete reaction, and determine the correct K value both in symbols (e.g., 1/Ka) and in number. • a) Potassium bicarbonate and hydrochloric acid • b) The acid dissociation of HCN • c) Acetic acid and potassium hydroxide

  34. So what really makes it strong vs. weak? • 1) Electronegativity: • Acid strength increases as electronegativity of A increases • 2) Bond strength: • Lessens as descend GVII  stronger acid • 3) Larger atomic radius: • Increases as descend G •  easier to lose H, stronger acid

  35. Continued • 3) Oxyacids • The more oxygens the stronger • Ex: in order of increasing acid strength • HClO < HClO2 < HClO3 < HClO4 • O-H bond polarity increases as oxygens are added  easier to remove hydrogen (as proton) • Inductive effect = ability of atoms in a molecule to attract electrons from another part of the molecule

  36. More… • Also, more oxygens allow for more stable (delocalized) structures • Compare Lewis structures of deprotonated (conjugate bases) HNO3 to HNO2 • Allows for greater acidity

  37. Carboxylic acids • Brønsted-Lowry acids • RCO2-H + H2O  RCO2- + H3O+ • Give me the resonance structures • If Electron Withdrawing Groups are substituted for hydrogens what would happen? • Order of increasing Ka • CH3CO2H > ClCH2CO2H > Cl2CHCO2H > Cl3CCO2H

  38. Anions as Brønsted bases • Basicity increased as negative charge of anion increases • PO43- • HPO42- • H2PO4- • Which is most basic? • Which has the highest pKb?

  39. Salts • Will a salt soln be acidic, basic, or neutral? • Let’s do these: HCl + NaOH  NaCl + H2O HCl + NH4OH  NH4Cl + H2O H3PO4 + 3CsOH  Cs3PO4 + 3H2O H3PO4 + 3NH4OH  (NH4)3PO4 + 3H2O

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