Honors Chemistry
Name: _____________________________________ Date: __________________ Mods: _________
The Common-Ion Effect and Buffered Solutions
17.1 The Common-Ion Effect:
 The extent of ionization of a weak electrolyte is ___________________ by adding to the solution a
_______________ electrolyte that has an ion in common with the weak electrolyte.
 Consider adding sodium acetate ( _______________) to a solution of acetic acid
(______________).
► Acetic acid is a weak acid which means it is a ___________ electrolyte.
► Sodium acetate is a soluble ionic compound according to our solubility rules which means it is
a ______________ electrolyte that dissociates ___________________ in aqueous solution
as seen below:
NaC2H3O2 (aq)  Na+ (aq) + C2H3O2- (aq)
 In contrast, HC2H3O2 ionizes as follows:
HC2H3O2 (aq)  H+ (aq) + C2H3O2- (aq)
 According to ______ ____________________ principle, the addition of C2H3O2-, from NaC2H3O2,
causes this equilibrium to shift to the __________, thereby __________________ the equilibrium
concentration of H+ (aq) and __________________ the pH.
1)
Example Problem: Calculating the pH When a Common Ion Is Involved
What is the pH of a solution made by adding 0.30 M acetic acid (HC2H3O2) and 0.25 M sodium
acetate (NaC2H3O2)? Ka for HC2H3O2 = 1.8 x 10-5
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*In Ch. 16, we calculated that a 0.30 M solution of HC2H3O2 has a pH of 2.64, corresponding to [H+] =
2.3 x 10-3 M. Thus, the addition of NaC2H3O2 has substantially decreased [H+] and raised the pH.
Now You Try!
Calculate the pH of a solution containing 0.085 M nitrous acid (HNO2; Ka = 4.5 x 10-4) and 0.10 M
potassium nitrite (KNO2).
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2)
Example Problem: Calculating Ion Concentrations When a Common Ion Is Involved
Calculate the fluoride ion concentration and pH of a solution that is 0.20 M in HF and 0.10 M in HCl.
Ka for HF = 6.8 x 10-4
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Now You Try!
Calculate the formate ion concentration and pH of a solution that is 0.050 M in formic acid (HCHO2;
Ka = 1.8 x 10-4) and 0.10 M in HNO3.
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 The ionization of a weak base is also ________________ by the addition of a common ion. For
example, the addition of NH4+ (as from the strong electrolyte, NH4Cl) causes the base-dissociation
equilibrium of NH3 to shift to the __________, _______________ the equilibrium concentration of
OH- and _______________ the pH as seen below:
NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq)
17.2 Buffered Solutions:
 Buffers are solutions of a _____________ conjugate acid-base pair.
 They are particularly resistant to changes in _______, even when strong acid or base is added.
► This is due to the fact that buffers have a _______________ component which can neutralize
H+ (from added acid) AND an _______________ component which can neutralize OH- (from
added base)
 To make a buffer: Add a ______________ acid or base with a _____________ of that acid or base
(recall that salts are ______________ compounds!)
► HC2H3O2- C2H3O2- buffer:
► weak acid- conj. base buffer
► add NaC2H3O2 (salt) to a solution of HC2H3O2 (weak acid)
► NH3 - NH4+ buffer:
► Weak base – conj. acid buffer
► add NH4Cl (salt) to a solution of NH3 (weak base)
Addition of an Acid or Base to a Buffer:
 If a small amount of base is added to a buffer of HF and F-, for example, the HF component of the
buffer will react with the added OH- , _____________________ [HF] and
_____________________ [F-] in the buffer.
 If a small amount of acid is added, the F- component of the buffer reacts with the added H+,
_____________________ [F-] and _____________________ [HF] in the buffer.
 To Summarize:
► Adding a base (OH-) to a buffer increases the concentration of conjugate base in a buffer
(very small __________________ in pH)
► Adding an acid (H+) to a buffer increases the concentration of the weak acid (very small
___________________ in pH)
 Buffer Calculations:
 Give the formula for the Henderson-Hasselbalch (H-H) equation:
► When [base] = [acid], ___________ = ___________
1)
Example Problem: Using the Henderson-Hasselbalch equation to find pH
What is the pH of a buffer that is 0.12 M in lactic acid, HC3H5O3, and 0.10 M in sodium lactate? Ka for
lactic acid is 1.4 x 10-4.
Now You Try!
Calculate the pH of a buffer composed of 0.10 M benzoic acid and 0.20 M sodium benzoate. Ka for
benzoic acid is 6.3 x 10-5.
2)
Example Problem: Using the Henderson-Hasselbalch equation to find a concentration
Calculate the concentration of lactic acid that must be present in a 0.014 M solution of sodium lactate
(NaC3H5O3) to produce a pH of 3.96. Ka for lactic acid is 1.4 x 10-4.
 Important Characteristics of Buffers:
1) Buffer capacity - the amount of acid or base the buffer can __________________ before the
pH begins to change to an ___________________ degree
2) pH Range – the range of pH values over which the buffer acts ____________________
► It is best to choose an acid with a pKa close to the desired pH.
Homework Practice Problems:
Directions: Use ICE table to solve the following common-ion solution problems.
1.
Calculate the pH of a solution that is 0.085 M propionic acid (HC3H5O2; Ka = 1.3 x 10-5) and
0.060 M potassium propionate (KC3H5O2).
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2.
Calculate the pH of a solution that is 0.260 M formic acid (HCHO2; Ka = 1.8 x 10-4) and 0.160 M
sodium formate (NaCHO2).
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Directions: Use the H-H equation to solve the following problems:
3.
Calculate the pH of a buffer that is 0.12 M in butanoic acid (HC4H7O2; Ka = 1.5 x 10-5) and 0.11 M
in lithium butyrate (LiC4H7O2).
4.
Calculate the Ka of a buffer that is 0.25 M in nitrous acid (HNO2) and 0.31 M in potassium nitrite
(KNO2) if the pH is 3.44.
5.
Calculate the concentration of sodium cyanate (NaCNO) that must be present in a 0.70 M
solution of cyanic acid (HCNO; Ka = 3.5 x 10-4) to produce a pH of 4.20.
6.
Calculate the concentration of chlorous acid (HClO2; Ka = 1.1 x 10-2) that must be present in a
0.005 M solution of sodium chlorite (NaClO2) to produce a pH of 2.45.