1 / 37

IV. Percent Ionization

IV. Percent Ionization. Instead of characterizing a weak acid by its K a , we can calculate how much it ionizes. The concentration of ionized acid is simply equal to the [H + ] at equilibrium. IV. Sample Problem. Find the % ionization of a 0.200 M acetic acid solution if its pK a = 4.74.

kieu
Télécharger la présentation

IV. Percent Ionization

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. IV. Percent Ionization • Instead of characterizing a weak acid by its Ka, we can calculate how much it ionizes. • The concentration of ionized acid is simply equal to the [H+] at equilibrium.

  2. IV. Sample Problem • Find the % ionization of a 0.200 M acetic acid solution if its pKa = 4.74.

  3. IV. Sample Problem • In a 0.0100 M solution of butyric acid at 20 °C, the acid is 4.0% ionized. Calculate the Ka and pKa of butyric acid at these conditions.

  4. IV. Mixtures of Acids • If there are multiple acids in solution, then there are multiple sources of H3O+. • If one of the acids is strong, it will be the major contributor, so much so that we can ignore contributions from others. HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) HCOOH(aq) + H2O(l)  H3O+(aq) + HCOO-(aq)

  5. IV. Mixtures of Weak Acids • If we have a mixture of weak acids, we need to examine the Ka’s. • If the Ka’s differ by more than a factor of several hundred, then we can just calculate based on the strongest acid. • If the Ka’s are too close together, we must solve two weak acid equilibria! • Start with the strongest weak acid and use that result in the second weak acid equilibrium.

  6. IV. Sample Problem • Calculate the pH of a solution that is 0.100 M in acetic acid, CH3COOH (Ka = 1.8 x 10-5), and 0.200 M in formic acid, HCOOH (Ka = 1.8 x 10-4).

  7. V. Basic Solutions • Just like acidic solutions, there are strong bases and weak bases. • Everything we learned about weak acids applies to weak bases. • The two systems are analogous to each other.

  8. V. Strong Bases • Strong bases ionize completely. • Strong bases are typically ionic compounds containing the hydroxide anion.

  9. V. Weak Bases • Weak bases typically do not produce OH- by partially ionizing. • Weak bases produce OH- by pulling a proton off water. B(aq) + H2O(l) BH+(aq) + OH-(aq)

  10. V. Weak Bases • The strength of a weak base depends on its base ionization constant, Kb. • For the generic weak base B(aq) + H2O(l) BH+(aq) + OH-(aq):

  11. V. Common Weak Bases

  12. V. Weak Base Structures • Weak bases tend to have lone pair e-’s that can accept a proton.

  13. V. Weak Base Problems • The method of solving weak base problems is no different than the method of solving weak acid problems! • Instead of Ka, you use Kb. • To get to pH, remember that you can go through pOH.

  14. V. Sample Problem • The pain reliever morphine is a weak base. In a 0.010 M morphine solution, the pH is 10.10. Calculate the Kb and pKb or morphine.

  15. V. Sample Problem • The pKb for pyridine is 8.77. What’s the pH of a 0.010 M aqueous solution of pyridine?

  16. VI. Ions as Weak Acids/Bases • Some ions can act as either weak acids or weak bases. • e.g. NH4+ and CH3COO- • These ions must be introduced into a solution as a salt. • e.g. NH4Cl and CH3COONa • These ionic salts ionize, and then the weak acid/base sets up its equilibrium.

  17. VI. Anions as Weak Bases • Any anion can be thought of as the conjugate base of an acid. • Anions that are conjugate bases of weak acids are themselves weak bases. • Anions that are conjugate bases of strong acids are pH neutral.

  18. VI. Conjugate Ka/Kb Pairs

  19. VI. Sample Problem • What’s the pH of a 0.10 M NaNO2 solution? Note that Ka for nitrous acid is 7.1 x 10-4.

  20. VI. Cations as Weak Acids • When cations go into aqueous solutions,we need to examine whether or not they will set up an equilibrium. • Cations of strong bases do nothing and are thus pH neutral. • Cations that are conjugate acids of weak bases will establish an acid equilibrium. • Small, highly-charged metal cations form weakly acidic solutions.

  21. VI. The Case of Al3+ • Al3+(aq) will form Al(H2O)63+ which will establish an acid equilibrium.

  22. VI. pH of Salt Solutions • To determine whether a salt solution will be acidic, basic, or neutral, we need to consider the nature of the cation and anion. • There are 4 possibilities: • Neither cation nor anion acts as acid or base. • Cation acts as acid, anion is neutral. • Anion acts as base, cation is neutral. • Cation acts as acid, and anion acts as base.

  23. VI. pH Neutral Salt Solutions • Cation comes from a strong base. • Anion comes from a strong acid.

  24. VI. Acidic Salt Solutions • Cation is conjugate acid of weak base or is a small, highly-charged cation. • Anion comes from a strong acid.

  25. VI. Basic Salt Solutions • Cation comes from a strong base. • Anion is a conjugate base of a weak acid.

  26. VI. “It Depends” • Sometimes, cation will be a conjugate acid and the anion will be a conjugate base. • In this case, the pH of the salt solution depends on relative values of Ka and Kb. • If Ka > Kb, solution will be acidic; if Kb > Ka, solution will be basic.

  27. VI. Analyzing Salt Solutions • Use the following steps to determine whether an aqueous solution of a salt will be acidic, basic, or neutral. • Break up salt into its cation and anion. • Ask yourself whether the cation can donate a proton or whether it is small and highly charged. If so, it is a weak acid. • Ask yourself whether the anion can accept a proton. If so, it is a weak base. • Consider the combined effect of having the cation and anion in solution.

  28. VI. Sample Problem • For each compound, predict whether its 0.1 M solution in water will be acidic, basic, or neutral. • NaNO2 • KCl • NH4Br • Fe(NO3)3 • NH4CN

  29. VII. Acids w/ More than one H+ • Some acids have more than one acidic proton; these are called polyprotic acids. • Generally, the Ka of the 2nd proton is much smaller than the 1st, so we generally just solve for the 1st ionization. • Exceptions: H2SO4 and when Ka’s are within a few hundred of each other. • For exceptions, it’s a double equil. problem!

  30. VII. Some Polyprotic Acids

  31. VII. Sample Problem • What is the pH and [SO42-] of a 0.0075 M sulfuric acid solution if Ka2 = 0.012 for sulfuric acid?

  32. VIII. Acid Strength & Structure • Why are some acids strong and some acids weak? • Depends on structure and composition of the acid. • We examine the factors that contribute to acid strength in binary acids and oxyacids.

  33. VIII. Binary Acids • The strength of a binary acid depends on bond polarity and bond strength. • The H must have the partial positive charge. • Weaker bond leads to greater acidity.

  34. VIII. Binary Group 16/17 Acids • The combined influence of polarity and bond strength can be seen in the Group 16 and Group 17 binary acids.

  35. VIII. Oxyacids • An oxyacid (a.k.a. oxoacid) has the general form H-O-Y- in which Y is some atom which may or may not have additional atoms bonded to it. • Oxyacid strength depends on the electronegativity of Y and the number of O atoms attached to Y.

  36. VIII. Electronegativity of Y • The more electronegative Y is, the more polar and weaker the O-H bond becomes. • If the O kicks off the H as H+, it can claim both electrons in the bond.

  37. VIII. # of O Atoms on Y • More O atoms draw electron density away from Y, which draws electron density from the O-H bond, leading to greater acidity.

More Related