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Chapter 16

Chapter 16. Acid-Base Equilibria. constant. Dissociation of water. Autoionization or autoprotolysis. Ion-product constant Autoprotolysis constant. K w = [H + ][OH - ] = 1.0x10 -14 When [H + ] = [OH - ] neutral. Doesn’t usually happen.

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Chapter 16

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  1. Chapter 16 Acid-Base Equilibria

  2. constant Dissociation of water Autoionization or autoprotolysis Ion-product constant Autoprotolysis constant

  3. Kw = [H+][OH-] = 1.0x10-14 When [H+] = [OH-] neutral. Doesn’t usually happen. As one increases, the other decreases; the product must equal 1.0x10-14. When [H+] > [OH-] acidic [OH-] > [H+] basic

  4. H+ is a proton with no electrons. In water: + é ù H O H ê ú H ë û Hydronium ion

  5. Bronstead-Lowry Acid-Base Acid - Can donate a proton Base - Can accept a proton *Doesn’t have to be in H2O. Can be in other solvents.

  6. conj base conj acid conj acid conj base Conjugate Acid-Base Pairs

  7. The stronger an acid, the weaker its conjugate base. The weaker an acid, the stronger its conjugate base.

  8. pH scale pH = -log [H+] Remember Kw = (1x10-7)(1x10-7) = 1.0x10-14 pH = -log [H+] = -log (1x10-7) pH = 7 (neutral) [H+] pH acidic > 1.0x10-7 < 7.00 basic < 1.0x10-7 > 7.00

  9. You can also speak in terms of [OH-] pOH = -log [OH-] = 14 - pH Because pH + pOH = -log Kw = 14

  10. Measure pH by pH meter Acid-base indicators Litmus red = pH < 5 blue = pH > 8 Figure 16.7 shows several acid-base indicators and their ranges

  11. Strong Acids and Bases Strong electrolytes Completely ionize HA + H2O  A- + H3O+ Bases form hydroxides in solvent

  12. In H2O, Alkali metal hydroxides Alkaline earth metal Hydroxides (except Be) Many are insoluble Also, substances that will abstract a H+ from H2O. O2- + H2O  2OH- Na2O or CaO would do this. O2-, H-, N3- bases that would do this.

  13. Weak acids Only partially ionize Acid dissociation constant

  14. O = C - O - H O N Larger Ka means stronger acid. ex. 0.020M solution pH = 3.26 ? Ka pH = -log [H+] = 3.26 [H+] = 5.50x10-4

  15. O O = = C - OH C - O O O N N  + H+ A- H+ HA 1:1

  16. O O = = C - OH C - O O O N N Can calculate pH in same manner if you have Ka and concentration of solution. Let’s use niacin again.  + H+ A- H+ HA

  17. ** Simplifying Assumption ** x is very very small compared to 0.010M sooooooooo, ignore x in denominator

  18. pH = -log [H+] x = [H+] = 3.9x10-4 pH = 3.41 What percent of niacin molecules ionized?

  19. Polyprotic Acids ex. H2SO4 H3PO4 H2SeO4 H2SO4 H+ + HSO4- Ka1 = 1.7x10-2 HSO4- H+ + SO42- Ka2 = 6.4x10-8 Ka1 always larger than Ka2 If Ka1/ Ka2 103, can estimate pH by Ka1 only.

  20. Weak Bases ex. Amines “an organic substituted ammonia” methyl amine ammonia NH3 CH3 H N H N H H H

  21. H CH3 + H2O  H N H N CH3 + OH- H H Anions of weak acids ClO- + H2O  HClO + OH- Kb = 3.3x10-7 Can use this in the same manner in which you used Ka.

  22. Ka and Kb How are they related?

  23. 1) 2) 3) When two reactions are added together, the equilibrium constant for the third reaction is given by the product of equilibrium constants of equations 1 and 2.

  24. K1 x K2 = K3 rxn 1 rxn 2 rxn 3

  25. Special Case Ka x Kb = Kw For conjugate acid-base pairs.

  26. Bond polarity and Bond strength effect on Acid-base behavior: In binary acids  polarity(across a row)  acidity  bond strength(in a group) acidity  stability of conj. base  acidity

  27. Metal hydrides are basic or show no acid/base properties in H2O. Nonmetal hydrides are acidic or show no acid/base properties in H2O (except NH3) Acidity increases moving down a group.

  28. Oxyacids O Have unprotonated and protonated oxygens. H O S O H O H3PO4 Y O H • As electronegativity of Y increases, acidity increases. • As number of unprotonated oxygens increases, acidity increases (effect of formal charge and oxidation number) • Ex. HClO, HClO2, HClO3, HClO4

  29. Carboxylic Acids O COOH = Carboxyl group C R OH R = H or an organic group. The more electron withdrawing R is, the greater the acidity (this stabilizes anion and weakens O-H bond) ex. H F O O H C C F C C O O H H F H Acetic acid Ka = 1.8x10-5 Trifluoroacetic acid Ka = 5.0x10-1

  30. Lewis Acids and Bases This is a completely different definition for acid/base chemistry than what you have seen thus far!!! Lewis acid = electron pair acceptor Lewis base = electron pair ‘donor’ Not giving them away, just has them available to ‘share’.

  31. H+ Bronstead-Lowry acid also a Lewis acid H+ electron pair acceptor OH- Electron pair donor Lewis base also Bronstead-Lowry base

  32. H H BH3 not a Bronstead-Lowry acid, but it’s a Lewis acid B H Incomplete Octet H Lewis Base has an electron pair available to attack an area that is e- deficient N H H

  33. Transition metal ions are often Lewis Acids. They have vacantdorbitals. (s and p also) H H O O = C = O Can be a Lewis Acid because e- density around the C is bound in just 2 directions.

  34. O H H O H C = = H O O H O O H C = = O C = = O O Carbonic acid Hydrolysis of metal ions Metal ions have positive charge so they attract the lone e- pair on H2O molecules

  35. 3+ H H H O O O Fe3+ H H H Fe H H O O O O H H H H H H 6 of these Because the metal is (+), e- density of H2O moves toward the metal. When this happens, there is less e- density in water’s O-H bonds, so H+ can come off easier…  pH will drop.

  36. The higher the charge density of the metal ion, the greater the acidity of its aqua complex.

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