1. Applications of Aqueous Equilibria
Chapter 15
Applications of Aqueous
Equilibria

2. Common Ion Effect HF(aq) H+(aq) + F-(aq) NH3(aq) NH4+(aq) + OH-(aq)
When a salt with the anion of a weak acid is added to that
acid (or the cation of a weak base is added to that base), it
reverses the dissociation.
HF(aq) H+(aq) + F-(aq)
NH3(aq) NH4+(aq) + OH-(aq)

3. What is a Buffer? A weak acid and its conjugate base is present
in some form (usually within a salt) OR
A weak base and its conjugate acid is present

4. pH of Buffers Practice Calculate the pH of a solution containing a
mixture of M HC3H5O2
(Ka = 1.3 x 10-5) and M NaC3H5O2
Calculate the pH of a solution containing a
mixture of 0.25 M NH3 (Kb = 1.8 x 10-5) and
0.10 M NH4Br.

5. Henderson-Hasselbalch Equation
Useful for calculating pH when the [A-]/[HA]
Ratios are known.
pH = pKa + log [A-]/[HA]
pH = pKa + log [base]/[acid]
Only can be used if you are dealing with a buffer system!!!
[H+] pH
Ka pKa

6. Calculate the pH of the following: 0.75 M lactic Acid (HC3H5O3, Ka = 1.4 x 10-4) with 0.25 M sodium lactate (NaC3H5O3)

7. A Buffered Solution
…….resists change in its pH when either H+ or OH- are added. 1.0 L of 0.50 M H3CCOOH M H3CCOONa pH = 4.74
Adding mol solid NaOH raises the pH of
the solution to 4.76, a very minor change

8. Adding a Strong Acid…….
….to a weak acid buffer system:

9. Adding a Strong Acid……. ….to a weak acid buffer system:
A strong acid will add its proton to the
conjugate base
…..to a weak base buffer system:

10. Adding a Strong Base…….
….to a weak acid buffer system:

11. Adding a Strong Base……. ….to a weak acid buffer system:
A strong base will add grab protons from
the weak acid
…..to a weak base buffer system:

12. Buffer Practice
Write the reaction that occurs when NaOH is added to the following buffer solutions: 1.00 M HNO2/1.00M NaNO2 Write the reaction that occurs when HCl is added to the following buffer solutions: M NH3/0.15 M NH4Cl

13. Buffering Capacity
represents the amount of H+ or OH- the buffer can absorb without a significant change in pH
The pH of a buffered solution is determined by
the ratio [A-]/[HA]
The capacity of a buffered solution is determined
by the magnitudes of [A-]/[HA]

14. Buffering Capacity Practice
Calculate the pH of each buffer system below: 5.00 M HAc and 5.00 M NaAc 0.05 M Hac and M NaAc Ka = 1.8 x 10-5 After HCl(g) is added to each buffer system, the first one has a resulting pH of 4.74, and the second one has a resulting pH of Which one has the better buffering capacity?

15. Review Session #2
Are You Ready?

16. How to Choose a Buffer The most effective buffers (most resistant to
change) have a ratio [A-]/[HA] = 1
True when [A-] = [HA]
Make pH = pKa (since log 1 = 0)
When choosing a buffer for a desired pH, choose a
buffer system whose pKa is close to the desired pH.

17. Review Session #3
Spring Break

18. TITRATION CURVES
We can determine the amount of a certain substance by performing
a technique called “titration.”
Analyte

19. for a Successful Titration:
3 Criteria Must Be Met
for a Successful Titration:
Exact reaction between titrant and analyte
MUST be known.
Equivalence point MUST be marked accurately.
Volume of titrant required to reach the
equivalence point MUST be accurate.
Note: When analyte is an acid or base…titrant is a
strong base or acid, respectively.

20. Titration (pH) Curve A plot of the solution being analyzed as a
function of the amount of titrant added.
Equivalence
(stoichiometric) point:
Enough titrant has been
added to react exactly
with the solution being
analyzed.
Moles acid = moles of base

21. Strong Acid-Strong Base Titration
WMX : W = Water; M = Molarity; X = add or sub. NmN : N = Neutralization; m = moles; N = numbers

22. Strong Acid-Strong Base Titration
They both dissociate completely
Do the stoichiometry
There is no equilibrium
The titration of 40.0 mL of M HCl04
with M KOH
Calculate the pH after the following volumes of
KOH have been added: 0,10.0, 40.0, 80.0 and
100.0 mL

23. Strong Acid-Strong Base Titration
The titration of 40.0 mL of M HCl04 with M KOH
Calculate the pH after the following volumes of
KOH have been added: 0,10.0, 40.0, 80.0 and
100.0 mL

24. Strong Acid-Base Titration

25. Weak Acid-Strong Base Titration
Consider the titration of mL of M acetic acid (Ka = 1.8 x 10-5) by M KOH. Calculate the pH of the resulting solution after The following volumes of KOH have been added: 0,50.0, 100.0, and mL

26. Weak Acid-Strong Base Titration
Consider the titration of mL of M acetic acid (Ka = 1.8 x 10-5) by M KOH. Calculate the pH of the resulting solution after The following volumes of KOH have been added: 0,50.0, 100.0, and mL

27. Titration Practice A beaker contains 100.0 mL of a solution of
HOCl (Ka = 3.5 x 10-8) of unknown concentration.
The solution was titrated with M NaOH, and the equivalence point was reached when 40.0 mL of NaOH was added. What was the original (initial) concentration of the HOCl solution?
What is the concentration of OCl- ions at the equivalence point?
What is the pH of the solution at the equivalence point?

28. Weak Acid-Strong Base Titration
The pH at the
equivalence point
of a titration of a
weak acid with a
strong base is
always greater
than 7

29. Weak Base-Strong Acid Titration Summary
Consider the titration of mL of M NH3 (Kb = 1.8 x 10-5) by M HCl. Calculate the pH of the resulting solution after the following volumes of HCl have been added: 0, 20.0, 25.0, 50.0 and mL

30. Weak Base-Strong Acid Titration
The pH at the
equivalence point
of a titration of a
weak base with a
strong acid is
always less
than 7

31. Acid-Base Indicators
Used in titrations to indicate when we have passed the equivalence point.
Weak acids that change color when they become bases
HIn H In-
Red
yellow
Equilibrium is controlled by pH
End point – when the indicator changes
color….tells us we have reached the eq. point.

32. Phenolphthalein Ka = 1.9 x 10-9
Acid-Base Indicators
Phenolphthalein Ka = 1.9 x 10-9
What is the pKa? 9
Bottom Line…….Take Home Message
pH of color change = pKa +/- 1

33. Solubility Product For solids dissolving to from aqueous solutions:
Bi2S3(s) Bi3+(aq) S2-(aq)
Ksp = solubility product constant
and
Ksp = [Bi3+]2 [S2-]3

34. Solubility Product
“Solubility” = S = concentration of Bi2S3 that dissolves , which equals ½[Bi3+] and 1/3[S2-].
Note: Ksp is constant (at a given temperature)
S is variable (especially with a common ion present)

35. Solubility Product Practice
The solubility of BiI3 is 1.32 x 10-5 mol/L.
Calculate the Ksp value.
The solubility of Ca3(PO4)2 is 2.05 x 10-7 mol/L
Calculate the Ksp value
The Ksp value for PbBr2 is 4.6 x 10-6
Calculate the solubility at 25oC.
The Ksp value for Ag2CO3 is 8.1 x 10-12

36. Relative Solubilities
Ksp will only allow us to compare the solubility of solids that fall apart into the same number of ions.
The bigger the Ksp of those the more soluble
PbBr2
Ag2CO3
If they fall apart into different number of pieces you have to do the math! (see #3 & 4) from previous slide)

37. Common Ion Effect If we try to dissolve the solid in a solution
with either the cation or anion already present, less solid will dissolve.
1. Calculate the solubility of solid SrSO4 (Ksp = 3.2 x 10-7) in a solution of 0.01M Na2SO4.
2. Calculate the solubility of solid Ca3(PO4)2 (Ksp = 1.3 x 10-32) in a solution of 0.20M Na3PO4.

38. pH and Solubility If the anion X- is an effective base – that is,
HX is a weak acid – salt MX will show
increased solubility in an acidic solution.
CaC2O Ca2+ + C2O42-
H+ + C2O HC2O4-