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Acid-Base Equilibria

Acid-Base Equilibria. Sour taste React with active metals to release hydrogen gas Change the color of indicators. Bitter taste Feel slippery Change the color of indicators. Acids Bases. Arrhenius Acid/Base Theory.

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Acid-Base Equilibria

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  1. Acid-Base Equilibria

  2. Sour taste React with active metals to release hydrogen gas Change the color of indicators Bitter taste Feel slippery Change the color of indicators Acids Bases

  3. Arrhenius Acid/Base Theory • According to Arrhenius, acids are substances that produce H+ (H3O+) ions in water. • Bases are substances that produce OH- ions in water.

  4. Why does water affect acids? • A water solution of HCl produces: HCl(g)  H+(aq) + Cl-(aq) • Acidic solutions are formed when an acid transfers a proton to water. H+ + H2O  H3O+

  5. Brønsted-Lowry Acids/Bases • Acids are substances capable of donating a proton. • Bases are capable of accepting a proton. • B/L theory applicable to reactions that do not occur in water • Can include gas phase reactions • NH3(g) + H2O(l)  NH4+(aq) + OH-(aq)

  6. An acid and a base which differ by only one proton are conjugate acid/base pairs.

  7. The more readily an acid gives up a proton, the less readily the conjugate base will accept a proton. The stronger the acid, the weaker its conjugate base and the weaker the acid, the stronger its conjugate base. HCN is a weaker acid than HF. The conjugate base CN- is a stronger base than F-.

  8. Diethylammonium ion (CH3)2NH2+ is a weak acid. Determine the conjugate base: Is the conjugate base stronger or weaker than the conjugate base of HCl?

  9. Strong Acids • Ionize completely to form ions • HCl HBr HI HNO3 H2SO4 HClO4

  10. Autoionization of Water • Proton transfer from one water molecule to another H-O-H + H-O-H  H3O+ + OH- H2O  H + + OH- K = [H+][OH-] H2O Kw = [H+][OH-] = 1.0x10-14

  11. [H2O] >> 55 M and remains constant  not written in K If solution is: neutral [H+] = [OH-] acid [H+] > [OH-] base [H+] < [OH-] [H+] = 2x10-5M [OH-] = 3x10 –9M [OH-] = 1 x 10-7M

  12. pH Scale • pH = -log[H+] • Acidic solution pH < 7 • Basic solution pH > 7 • Neutral solution pH = 7 • Number of decimal places in pH = number of significant figures in concentration. 1.0x10-12 pH = 12.00

  13. Strong Acids • Strong electrolytes • React completely with water to form H+ • pH of a strong acid equals –log of hydrogen concentration

  14. Weak Acids • Weak electrolytes • [molecules] greatest at equilibrium • H atoms bound to C do not ionize while H atoms bound to O will ionize HX(aq)  H+(aq) + X-(aq) Ka = [H+][X-] [HX]

  15. 0.10M HCOOH (formic acid) pH = 2.38 Calculate Ka and percent dissociation

  16. In the reaction of a strong acid with metal the conductivity remains constant since all acid molecules ionize. But, for a weak acid, conductivity will fluctuate since degree of ionization increases As [A] decreases. Calculate %HF in 0.10 M HF 0.010 M HF

  17. Polyprotic Acids • Release more than one proton in H2O H2SO3 H+(aq) + HSO3-(aq) Ka1 HSO3-(aq)  H+(aq) + SO32-(aq) Ka2 Ka1 = 1.7 x 10-2 Ka2 = 6.4 x 10-8

  18. If Ka constants differ by a factor of 103, consider Ka1 only. In other words, treat the acid as if it were monoprotic. The solubility of CO2 at 25oC and 0.10 atm equals 0.0037 M. All of the dissolved CO2 is as H2CO3. What is the pH of a 0.0037 M solution of H2CO3?

  19. Strong Bases • Most common strong bases are heavy Group IIA and all Group IA. • Determine the pOH of a basic solution by: pOH = -log [OH-] pH + pOH = 14 Determine the pH of a 0.010 M solution of Ba(OH)2.

  20. Formation of Strong Bases • Stronger bases than water are able to remove an H+ ion from water: O2- + H2O(l)  2OH-(aq) H- + H2O(l)  H2(g) + OH-(aq) N3- + H2O(l)  NH3(aq) +3OH-(aq)

  21. Weak Bases • Weak base + water  acid + OH- NH3 + H2O  NH4+ + OH- Kb = [NH4+][OH-] [NH3]

  22. Calculate the [OH-] in a 0.15 M solution of NH3. (Kb = 1.8 x 10-5)

  23. Amines • Weak nitrogen bases • N-C bonds • Due to lone pair on N, it is able to extract a proton H2N-CH3 + H2O  [H3NCH3]+ + OH-

  24. Anions of Weak Acids When sodium salts dissolve in water, the Na+ ion merely acts as a spectator ion. The reaction occurs between the remaining anion and water. C2H3O2- + H2O  HC2H3O2 + OH- Calculate the pH of a 0.010 M solution of NaClO.

  25. Relation Between Ka and Kb NH4+(aq)  NH3(aq) + H+(aq) NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) NH4+ and NH3 are conjugate pairs. Ka = [NH3][H+] Kb = [NH4+][OH-] [NH4+] [NH3] Add Reaction 1 and Reaction 2.

  26. H2O  H+ + OH- Multiply Ka x Kb Ka x Kb = Kw Ka and Kb are inversely related, as acid strength increases (Ka), base strength decreases (Kb) since product must equal Kw.

  27. Calculate the Kb for F- and the Ka for NH4+. Kb = 1.5 x 10-11 Ka = 5.6 x 10-10 Which of the following has the largest Kb? NO2-, PO43-, N3- PO43-

  28. Acid-Base Properties of Salts • Completely ionized • Nearly all salts are strong electrolytes • Acid/base properties due to hydrolysis of cations and anions • Strong acids and bases produce ions that do not hydrolyze

  29. Anions of weak acids: NO2-(aq) + H2O(l)  HNO2(aq) + OH-(aq) • Cations of weak bases: NH4+(aq) + H2O(l)  NH3(aq) + H3O+(aq) • Anions of polyprotic acids act as either proton donors or acceptors depending on the value of Ka or Kb. Predict whether Na2HPO4 will be acidic or basic.

  30. pH of Salts • Depends on parent acid and base. • If acid and base are both strong, salt is neutral. • If acid is strong and base is weak, salt is acidic. • If acid is weak and base is strong, salt is basic. • If parent acid and base are both weak, salt pH depends on value of Ka and Kb. List the following in order of increasing pH: Co(ClO4)2 RbCN Sr(NO3)2 KC2H3O2

  31. Acid/Base Character • Any molecule containing H can act as a potential acid, but bond must be polarized. • Very strong bonds are less ionized than weak ones. HF is a weak acid but acidic character increases with increasing atomic number making the remaining elements in the halogen family strong acids.

  32. OxyAcids • Acids in which OH groups and possibly additional oxygen atoms are bound to a central atom. • Acid strength increases with increasing EN of Y. HIO < HBrO < HClO

  33. In a series of acids with the same Y, acid strength increases with increasing oxidation number. HClO < HClO2 < HClO3 < HClO4

  34. Lewis Acids/Bases • According to Lewis, an acid is an electron pair acceptor. Acids will generally have incomplete octets. • A base is an electron pair donor. Bases will tend to have lone pairs. BF3 + NH3 F3BNH3

  35. http://antoine.frostburg.edu/chem/ senese/101/acidbase/indicators.shtml

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