0
10
20
30
40
50
Buffers
and
Acid/Base
Titration
The Common Ion Effect
• When the salt with the anion of a weak
acid is added to that acid, it reverses the
dissociation of the acid.
• It lowers the percent dissociation of the
acid.
• This same principle applies to salts with
the cation of a weak base.
Buffered Solutions
A solution that resists a change in
pH when either hydroxide ions or
protons are added.
Buffered solutions contain either:
 A weak acid and its salt
 A weak base and its salt
We can make a buffer of any pH by
varying the concentrations of these
solutions.
Acid/Salt Buffering Pairs
The salt will contain the anion of the acid,
and the cation of a strong base (NaOH, KOH)
Weak Acid
Formula
of the acid
Hydrofluoric
HF
Formic
HCOOH
Benzoic
C6H5COOH
Acetic
Carbonic
Propanoic
Hydrocyanic
CH3COOH
H2CO3
HC3H5O2
HCN
Example of a salt of the
weak acid
KF – Potassium fluoride
KHCOO – Potassium formate
NaC6H5COO – Sodium benzoate
NaH3COO – Sodium acetate
NaHCO3 - Sodium bicarbonate
NaC3H5O2 - Sodium propanoate
KCN - potassium cyanide
Base/Salt Buffering Pairs
The salt will contain the cation of the base,
and the anion of a strong acid (HCl, HNO3)
Formula of
the base
Example of a salt of the weak
acid
NH3
NH4Cl - ammonium chloride
Methylamine
CH3NH2
CH3NH2Cl – methylammonium chloride
Ethylamine
C2H5NH2
C2H5NH3NO3 - ethylammonium nitrate
Aniline
C6H5NH2
C6H5NH3Cl – aniline hydrochloride
Base
Ammonia
Pyridine
C5H5N
C5H5NHCl – pyridine hydrochloride
Adding a Strong Acid or Base to
a Buffered Solution
• Do the stoichiometry first
–Use moles!
• Next do the equilibrium problem
–ICE, ICE baby!
Titration of an Unbuffered Solution
13
12
11
10
A solution that is
0.10 M CH3COOH
is titrated with
0.10 M NaOH
9
pH
8
7
6
5
4
3
2
1
0.00
5.00
10.00
15.00
20.00
25.00
milliliters NaOH (0.10 M)
30.00
35.00
40.00
45.00
Titration of a Buffered Solution
13
12
11
A solution that is
0.10 M CH3COOH and
0.10 M NaCH3COO is
titrated with
0.10 M NaOH
10
9
pH
8
7
6
5
4
3
2
1
0.00
5.00
10.00
15.00
20.00
25.00
milliliters NaOH (0.10 M)
30.00
35.00
40.00
45.00
Comparing Results
Gra ph
pH
Buffered
Unbuffered
mL 0.10 M NaOH
Comparing Results
Buffered
13
13
12
12
11
11
10
10
9
9
8
8
7
pH
pH
Unbuffered
6
6
5
5
4
4
3
3
2
1
0.00
7
2
5.00
10.00
15.00
20.00
25.00
milliliters NaOH (0.10 M)
30.00
35.00
40.00
45.00
1
0.00
5.00
10.00
15.00
20.00
25.00
30.00
35.00
milliliters NaOH (0.10 M)
In what ways are the graphs different?
In what ways are the graphs similar?
40.00
45.00
Henderson-Hasselbach Equation
 [ A ] 
 [base] 
  pK a  log 

pH  pK a  log 
 [acid ] 
 [ HA] 
 [ BH  ] 
 [acid ] 
  pKb  log 

pOH  pKb  log 
 [base] 
 [ B] 
This is an exceptionally powerful tool, and it’s use
will be emphasized in our problem solving. I
memorized and use the pH version of the H-H Eq.
Who loves Henderson-Hasselbach, I do!
Buffer Capacity
• The pH of a buffered solution is determined by
the ratio [A-]/[HA].
• As long as this doesn’t change much, the pH
won’t change much.
• The more concentrated these two are, the
more H+ and OH- the solution will be able to
absorb.
• Larger concentrations = bigger buffer capacity
Buffer Capacity
•
•
•
•
The best buffers have a ratio [A-]/[HA] = 1
This is most resistant to change
This is true when [A-] = [HA]
As a result pH = pKa (since log 1 = 0)
Titrations
• Millimole (mmol) = 1/1000 mL
• Molarity = mmol/mL = mol/L
• This makes calculations easier because
we will rarely add liters of solution.
• Adding a solution of known concentration
until the substance being tested is
consumed.
• This is called the equivalence point.
• Graph of pH vs. mL is a titration curve.
Strong Acid/Strong Base Titration
13
12
11
10
9
pH
8
7
Endpoint is at
pH 7
A solution that is
0.10 M HCl is
titrated with
0.10 M NaOH
6
5
4
3
2
1
0.00
5.00
10.00
15.00
20.00
25.00
milliliters NaOH (0.10 M)
30.00
35.00
40.00
45.00
Weak Acid/Strong Base Titration
13
12
11
10
9
Endpoint is above
pH 7
pH
8
7
A solution that is
0.10 M CH3COOH
is titrated with
0.10 M NaOH
6
5
4
3
2
1
0.00
5.00
10.00
15.00
20.00
25.00
milliliters NaOH (0.10 M)
30.00
35.00
40.00
45.00
Strong Acid/Strong Base Titration
13
12
A solution that is
0.10 M NaOH is
titrated with
0.10 M HCl
11
10
9
pH
8
7
Endpoint is at
pH 7
It is important to
recognize that
titration curves are
not always
increasing from left
to right.
6
5
4
3
2
1
0.00
5.00
10.00
15.00
20.00
25.00
milliliters HCl (0.10 M)
30.00
35.00
40.00
45.00
Strong Acid/Weak Base Titration
13
12
11
10
9
pH
8
7
6
5
Endpoint is below
pH 7
4
A solution that is
0.10 M HCl is
titrated with
0.10 M NH3
3
2
1
0.00
5.00
10.00
15.00
20.00
25.00
milliliters NH3 (0.10 M)
30.00
35.00
40.00
45.00
Strong Acid-Strong Base
Titration
• Do the stoichiometry
– mL x M = mmol
• Remember, there is no equilibrium
because they both dissociate completely.
• The reaction is H+ + OH-  HOH
Weak Acid-Strong Base
Titration
• Do the stoichiometry using moles
• Determine the major species
• Do the equilibrium
WA-SB Titration Summary
• Weak acid with 0 mL of base—use Ka
• Before the equivalence point
– Stoichiometry first
– Then Henderson-Hasselbach
• At the equivalence point—use Kb
• After the equivalence point you have
leftover strong base
Strong Acid-Weak Base
Titration
• Process is basically the same as Weak
Acid-Strong Base
• Before equivalence point do the stoich and
then use H-H eq.
• At the equivalence point use Ka
• After the equivalence point you have left
over strong acid.
Indicators
• They are weak acids that change color
when they become bases.
• Written as Hin
• Ex: HIn  H+ + Inclear
red
• The equilibrium is controlled by pH
• End point = when the indicator changes
color.
• Try to match the equivalence point.
Indicators
• Since it is an equilibrium the color change is
gradual.
• It is noticeable when the ratio of
[In-]/[HI] or [HI]/[In-] is 1/10
• Since an indicator is a weak acid, it has a Ka
• The pH the indicator changes at is:
pH = pKa + log([In-]/[HI]) = pKa + log(1/10)
pH = pKa – 1 on the way up
Indicators
• pH = pKa + log([HI]/[In-]) = pKa + log(10)
• pH = pKa + 1 on the way down
• You want to chose an indicator with a pKa
1 more than the pH at the equivalence
point if you are titrating with a base.
• Choose the indicator with a pKa 1 less
than the pH at equivalence point if you are
titration with an acid.
Selection of Indicators
Indicator
Indicator Transitions
Low pH color
Transition pH range High pH color
Gentian violet (Methyl violet 10B)
yellow
0.0–2.0
blue-violet
Leucomalachite green (first transition)
Leucomalachite green (second transition)
yellow
green
0.0–2.0
11.6–14
green
colorless
Thymol blue (first transition)
red
1.2–2.8
yellow
Thymol blue (second transition)
yellow
8.0–9.6
blue
Methyl yellow
Bromophenol blue
red
yellow
2.9–4.0
3.0–4.6
yellow
purple
Congo red
blue-violet
3.0–5.0
red
Methyl orange
Bromocresol green
Methyl red
Methyl red
red
yellow
red
red
3.1–4.4
3.8–5.4
4.4–6.2
4.5–5.2
orange
blue
yellow
green
Azolitmin
Bromocresol purple
Bromothymol blue
Phenol red
Neutral red
Naphtholphthalein
Cresol Red
red
yellow
yellow
yellow
red
colorless to reddish
yellow
4.5–8.3
5.2–6.8
6.0–7.6
6.8–8.4
6.8–8.0
7.3–8.7
7.2–8.8
blue
purple
blue
red
yellow
greenish to blue
reddish-purple
Phenolphthalein
Thymolphthalein
Alizarine Yellow R
Litmus
colorless
colorless
yellow
red
8.3–10.0
9.3–10.5
10.2–12.0
4.5-8.3
fuchsia
blue
red
blue
Source: Wikipedia