CHAPTER 15: ACID -BASE EQUILIBRIA
15.1 SOLUTIONS OF ACIDS AND BASES CONTAINING A COMMON ION
Common Ion Effect : the addition of an ion already present in a system causes the equilibrium to shift __________
from the common ion.
Example: Suppose we have a hydrofluoric acid solution
Equilibrium Dissociation Equation:
How will the equilibrium shift with the addition of NaF?
How will the [H+] change?
pH change?
The extent of dissociation of HF will be ______________ in the presence of dissolved NaF.
Example Problem: Percent dissociation of HF in a 1.0 M solution is 2.7 %. Calculate the percent dissociation of HF in a
solution containing 1.0 M HF (Ka = 7.2 x 10-4) and 1.0 M NaF.
Major species?
Dissociation equation:
ICE Table
Write Ka expression, calculate [H+]
Calculate % Dissociation
15.2 BUFFERED SOLUTIONS
The most important application of acid-base solutions containing a common ion is for buffering.
A buffered solution is one that resists a change in ________________ when either ____ or ____
are added.
A buffered solution may contain a ___________ acid and its salt (ex
or a _____________ base and its salt (ex.
),
).
Example of a buffered solution: our blood – can absorb the acids and bases produced in biological reactions without
changing its pH (H2CO3/HCO3- buffer)
EXAMPLE 1: Calculate the pH of a solution that contains 0.250 M formic acid, HCOOH (Ka = 1.8 x 10-4) and 0.100 M
sodium formate, NaCOOH
Major species:
Dissociation equation and ICE table:
.
Ka expression and pH calculation:
Alternative method of pH calculation for a buffer:
The log form for the expression for Ka is called the Henderson-Hasselbalch equation
 [base] 

 [acid ] 
pH = pKa + log 
In doing equilibrium calculations, we have seen that we can normally neglect the amounts of the weak acid or weak base
of the buffer that ionize. Therefore, we can use the initial concentrations of the acid and base components of the buffer
directly in the above equation.
EXAMPLE 2: A buffered solution is made by adding 50.0 g NH4Cl to 1.00 L of 0.75 M solution of NH3. Calculate the
pH of the final solution (assume no volume change.) Kb for NH3 is 1.8 x 10-5.
pH CHANGES IN BUFFERED SOLUTIONS
Addition of a Strong Base to a Buffer
What happens when you add a strong base to a buffer?
Suppose a buffered solution contains relatively large quantities of weak acid, HA, and its conjugate base, A-.
When hydroxide ions are added to the solution, it will react with the weak acid HA:
This is a neutralization reaction that will go to completion!
The net result is that OH- ions are not allowed to
accumulate but are replaced by A- ions.
In a buffered solution, the pH is governed by the
ratio:
When OH- are added HA is converted to _____ ions, and the ratio
[HA]/[A-] _________________. If the amounts of HA and A- originally
present are very large compared with the amount of OH- added, the
change in [HA]/[A-] will be small.
For example: If 0.010 moles of NaOH is added to the formic
acid/formate ion buffer in previous Example 1(assume volume of 1L):
If 0.01 mol of NaOH had been added to 1.00 L of water instead, what would be the pH?
Similar reasoning applies when a strong acid is added to a buffered solution of a weak acid HA and its conjugate base,
A-. The H+ ions (from the strong acid) will react with the conjugate base, A- , again a neutralization reaction that goes to
completion:
Again, net change of A- to HA, however if [A-] and [HA] are large compared with the H+ added, little change to pH
results.
REVIEW:
A strong acid (H+) or a strong base (OH-) can also be added to a basic buffer system consisting of a weak base, B, and its
conjugate acid, BH+, with similar results:
If all the H+ or OH- added to a buffered solution are consumed, the buffer returns to an equilibrium state!
More Buffer Example Problems:
1) Calculate the mass of sodium acetate that must be added to 500.0 mL of 0.200 M acetic acid to form a
pH = 5.00 buffer solution. Ka for acetic acid is 1.8 x 10-5.
2) Consider a solution that contains both C5H5N and C5H5NHNO3. Calculate the ratio of [C5H5N]/[C5H5NH+]
if the solution has the following pH values:
(a) 4.50
(b) 5.23
15.3 BUFFER CAPACITY
The buffering capacity of a buffered solution represents the amount of hydrogen or hydroxide ions the buffer can absorb
without a significant change in pH. Once the OH- or H+ ions become the excess reactant in the neutralization reaction
involving the buffer system, the buffer capacity has been exceeded and the pH will change dramatically.
Example: A buffered solution contains 0.25 M NH3 (Kb = 1.8 x 10-5) and 0.40 M NH4Cl.
(a) Calculate the pH of this buffered solution.
(b) Calculate the pH that results when 0.50 mol of HCl is added to 1.0 L of the buffered solution from
part (a).
Therefore in making a buffer, we want quantities of acid and base in the buffer that give optimum or maximum
buffering.
* Optimum buffering will occur when
[base] = 1
[acid]
From the Henderson-Hasselbalch equation, pH =
and if [base]/[acid] = 1, then log(1) = ________
Therefore for optimum buffering pH = pKa
Preparing a Buffer:
Example 1: If we want to make a buffer with a pH of 5.55, we should choose an acid whose pKa
is closest to the desired pH of the buffer – in this case 5.55. Therefore we pick an acid whose Ka is closest to
______________.
Example 2: A chemist needs a solution buffered at a pH =4.30 and can choose from the following acids
(and their sodium salts). Which acid should the chemist choose?
a) hydrofluoric acid (Ka = 7.2 x 10-4)
b) nitrous acid (Ka = 4.0 x 10-4)
c) benzoic acid (Ka = 6.4 x 10-5)
HOMEWORK (separate sheet of paper)
Homework – 2/4
1. Calculate the % dissociation and pH for 0.100 M propanoic acid ,HC3H5O2 . (Ka = 1.3 x 10-5)
2. Calculate the pH of a mixture of 0.100 M propanoic acid and 0.100M NaC3H5O2.
3. Calculate the % dissociation of propanoic acid in the above mixture.
4. Calculate the pH of a buffered solution prepared by dissolving 21.5 g of benzoic acid (HC7H5O2) and
37.7 g sodium benzoate in 200.0 mL of solution. Ka for benzoic acid = 6.4 x 10-5
5. A certain buffer is made by dissolving NaHCO3 and Na2CO3 in some water. Write equations to show how
this buffer neutralizes added H+ and OH-. (Hint – remember the sodium ion is a spectator ion!)
Homework – 2/5
pp. 751-752 #18,20,33,36,42, and 44
Homework – 2/6
pp. 752 #46,48, and 52
Homework 2/7 – AP FRQ (will be given in class)
AP PROBLEM
Each of three beakers contains 25.0 mL of a 0.100 M solution of HCl, NH3 , or NH4Cl, as shown above.
Each solution is at 25°C.
(a) Determine the pH of the solution in beaker 1. Justify your answer.
(b) In beaker 2, the reaction NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq) occurs. The value of Kb for
NH3(aq) is 1.8 × 10−5 at 25°C.
(i) Write the Kb expression for the reaction of NH3(aq) with H2O(l).
(ii) Calculate the [OH-] in the solution in beaker 2.
(c) In beaker 3, the reaction NH4+(aq) + H2O(l) ⇌ NH3(aq) + H3O+(aq) occurs.
(i) Calculate the value of Ka for NH4+(aq) at 25°C.
(ii) The contents of beaker 2 are poured into beaker 3 and the resulting solution is stirred.
Assume that volumes are additive. Calculate the pH of the resulting solution.
(d) The contents of beaker 1 are poured into the solution made in part (c)(ii). The resulting solution is
stirred. Assume that volumes are additive.
(i) Is the resulting solution an effective buffer? Justify your answer.
(ii) Calculate the final [NH4+] in the resulting solution at 25°C.