EXPERIMENT 16
BUFFER SOLUTIONS
INTRODUCTION
A buffer is a solution that resists change in its pH upon addition of limited amounts of strong acid or
strong base or upon dilution.
Consider a buffer solution that is a mixture on the wea k acid, acetic acid, and its salt, sodium acetate. In
the buffer solution, the solute particles are:
C 2H3O2
Na +
HC2H3O2
Notice: the buffer solute particles include the weak acid HC2H3O2 and its conjugate base C2H3O2. These
are the particles that will do the buffering. The sodium ion is a spectator ion.
If a strong acid solution were added to the buffer solution, the C 2H3O2 ions in the buffer solution would
react with (neutralize) the H 3O+ ions from the strong acid as shown in the net ionic equation for the
neutralization reaction:
C2 H 3 O 2 
+
(from buffer solution)
H 3O+
(from strong
acid solution)
 HC2H3O2
+
H 2O
(1)
Or, if a strong base were added to the buffer solution, the HC 2H3O2 molecules in the buffer solution
would react with (neutralize) the OH  ions from the strong base as shown in the net ionic equation for
the neutralization reaction:
HC2H3O2
+
(from buffer solution)
OH 

(from strong
base solution)
C2H3O2
+
H 2O
(2)
The equations above show that a buffer solution (often called “buffer system”) will maintain its pH even
when considerable amounts of strong acid or base are added. Up until enough strong acid solution is
added to use up most of the C2H3O2 ions present in the buffer (see equat ion 1), or up until enough
strong base solution is added to use up most of the HC 2H3O2 in the buffer solution (see equation 2), the
buffer pH will remain constant to ±1.
The ability of a buffer solution to maintain its pH is called its buffering capacity. Greatest buffering
capacity is obtained if the buffer solution contains equimolar concentrations (equal molar concentrations)
of the weak acid and its conjugate base. In the case of our buffer system, that means
[HC2H3O2 ] = [C2H3O2]
In the buffer solution, HC 2H3O2 ionizes very little because of the presence of its conjugate base, C 2H3O2,
which is its common ion.
HC2H3O2 + H 2O

H3O+ + C2H3O2
therefore, at equilibrium, in the buffer solution, we can assume that
[HC2H3O2 ] = M HC2H3O2
Consequently,
[C 2H3O2]
[HC 2H3O2]
and
=
[C 2H3O2] = M C2H3O2
M C2 H 3 O 2 
M HC2H3O2
=
1
As a result, in a buffer solution containing equimolar concentrations of HC 2H3O2 and C 2H3O2, the
hydronium ion concentration, [H 3O+], is equal to the K a:
1
Ka = [H 3O+]
[C2H3O2]
[HC 2H3O2]
1
K a = [H3O+]
therefore:
and
Ka
[H 3O+]
=
[C2H3O2]
[HC 2H3O2]
HOW TO MAKE A BUFFER SOLUTION:
One way to make a buffer solution is to first select a weak acid that has a K a that is close to the
ACID
K
HS04
Ka = 1.1  10-2
Sulfurous acid
H2S03
Ka = 1.3  10-2
Phosphoric acid
H3P04
Ka = 6.9  10-3
Formic acid
HCH02
Ka = 1.7  10-4
Acetic acid
HC2H302
Ka = 1.7  10-5
Carbonic Acid
H2C03
Ka = 4.3  10-7
Hypochlorous acid
HClO
Ka = 3.5  10-8
Dihydrogen phosphate ion
H2P04
Ka = 6.2  10-8
Bicarbonate ion
HC03
Ka = 4.8  10-11
Ammonium ion
NH4+
Ka = 5.6 x 10 -10
BASE
Ammonia
Kb = 1.8  10-5
NH3
[H3O+] that you desire for your solution, or select a weak base whose K
Hydrogen sulfate ion
b
CONJUGATE BASE
S042
HS03
H2P04
CH02
C2 H 3 O 2 
HC03
ClO
HP042
C032
NH3
CONJUGATE ACID
NH4+
is close to the desired [OH ].
For example, if you wished to prepare a buffer solution of pH 4.20 —that is, [H3O+] = 6.3 x 10 -5, you would
look for the acid whose K a is closest to 6.3 x 10 -5.
In this case, the best choice is HC2H302 with Ka = 1.7  10-5, because it will give you the greatest buffering
capacity.
HC2H3O2 + H 2O  H3O+ + C2H3O2
You will then calculate the ratio of Ka to desired [H 3O+] which is equal to the ratio of conjugate base to acid
that you want to have in your buffer solution:
Ka
conjugate base
= [C 2H3O2]
+
[H 3O ]
[HC 2H3O2]
acid
1.7 x 10 -5 =
6.3 x 10-5
0.27 =
1
[C 2H3O2]
[HC 2H3O2]
Therefore the ratio of M C2H3O2 to M HC2H302 in the buffer solution is 0.27 to 1. And since we will be
using equimolar concentrations of conjugate base and acid, the ratio of the volume of conjugate base
solution to volume of acid solution in the buffer will be 0.27 to 1
M C2 H3 O2 
M HC2H302
=
0.27
1
=
V conjugate base soluton
V acid solution
2
To prepare a particular volume of this buffer, in which there are .27 parts conjugate base to 1 p art acid,
first add the parts to get the total parts (in this case 1 + .27 = 1.27), and then divide the buffer solution
volume by the total parts to get the volume per part. For example, if we wanted 30 mL of this buffer,
30 mL
1.27 parts
=
23.6 mL/part
Therefore the volume of acid solution is:
1 part acid x 23.6 mL/part = 23.6 mL acid solution
And the volume of conjugate base solution is: 30.0 mL – 23.6 mL = 6.4 mL conjugate base solution
The buffer solution is then prepared by adding 23.6 mL of the acid solution to 6.4 mL of the conjugate
base solution.
EXPERIMENT
Check out a pH pen from the stockroom
A.
AMMONIUM ACETATE BUFFER SOLUTION + STRONG ACID
1.
Measure the pH of 1.0 M NH4C2H3O2 solution. Record the pH on the report sheet.
2.
Prepare 1 M HCl solution:
Prepare about 25 mL of 1 M HCl solution by diluting 6 M HCl (bottle #2 at your bench) with
deionized water. MIX WELL.
Calculate volume of 6 M HCl needed:
3.
Label 2 test tubes and prepa re the following mixtures.
Tube #1:
Tube #2:
5 mL deionized water + 2 drops methyl orange indicator
5 mL 1.0 M NH4C2H3O2 (buffer) solution + 2 drops methyl orange indicator
4.
To Tube #1 add the 1 M HCl solution prepared in step 2 above drop by drop, stirring, and
counting the number of drops needed to change the color of methyl orange (at about pH 3).
5.
Measure 10 mL of the 1 M HCl solution into your 10 mL graduated cylinder. Add the 1 M HCl
solution to Tube #2 until the color of methyl orange has changed (about pH 3). You do not
need to count drops, instead, read the final volume of HCl solution in your graduated cylinder
and calculate the volume HCl add ed by taking the difference between 10 mL and the final
volume. Record the volume on your report sheet.
SAVE THE REMAINNG 1 M HCl SOLUTION FOR USE IN PART C
3
B.
AMMONIUM ACETATE BUFFER SOLUTION + STRONG BASE
1.
Prepare 1 M NaOH solution:
Prepare about 25 mL of 1 M NaOH solution by diluting 6 M NaOH (bottle #7 at your bench) with
deionized water. MIX WELL.
Calculate volume of 6 M NaOH needed:
2. Label 2 test tubes and prepare the following mixtures.
Tube #1:
Tube #2:
5 mL deionized water + 2 dr ops Alizarin Yellow R indicator
5 mL 1.0 M NH4C2H3O2 (buffer) solution + 2 drops Alizarin Yellow R indicator
3. To Tube #1 add the 1 M NaOH solution prepared in step 2 above drop by drop, stirring, and counting
the number of drops needed to change the color of Alizarin Yellow R (at about pH 11).
4. Measure 10 mL of the 1 M NaOH solution into your 10 mL graduated cylinder. Add the 1 M
NaOH solution to Tube #2 until the color of Alizarin Yellow R has changed (about pH 11). You
do not need to count drops, instead, read the final volume of NaOH solution in your graduated
cylinder and calculate the volume NaOH added by taking the difference between 10 mL and the
final volume. Record the volume on your report sheet.
SAVE THE REMAINNG 1 M NaOH SOLUTION FOR USE IN PART C
C.
SODIUM BICARBONATE BUFFER
1.
Measure the pH of 0.10 M NaHCO 3 solution. Record the pH on the report sheet.
2.
Label 2 test tubes and prepare the following mixtures.
Tube #1:
Tube #2:
5.00 mL 0.10 M NaHCO 3 + 2 drops methyl orange indicator
5.00 mL 0.10 M NaHCO 3 + 2 drops Alizarin Yellow R indicator
3. Measure 10 mL of the 1 M HCl solution into your 10 mL graduated cylinder. Add t he 1 M HCl
solution to Tube #1 until the color of methyl orange has changed (about pH 3). Record the
volume 1 M HCl solution required on your report sheet.
4.
Measure 10 mL of the 1 M NaOH solution into your 10 mL graduated cylinder. Add the 1 M
NaOH solution to Tube #2 until the color of Alizarin Yellow R has changed (about pH 11).
Record the volume of 1 M NaOH solution required on your report sheet.
4
D. PREPARATION OF A BUFFER SOLUTION:
Your instructor will assign you a pH for your buffer solution. You will then prepare 20 to 30 mL of
the buffer solution To do this, first choose from the table below the acid whose K a is closest to
your assigned [H3O+].
Acid
CONJUGATE
BASE
Ka
SOLUTIONS FOR BUFFER
PREPARATION
ACID OR
SALT OF
ACID SALT
CONJUGATE BASE
hydrogen sulfate ion
HS04
Ka = 1.1  10-2
S042
0.1 M NaHS04
0.1 M Na2S04
formic acid
HCH02
Ka = 1.7  10-4
CH02
0.1 M HCH02
0.1 M NaCHO 2
acetic acid
HC2H302
Ka = 1.7  10-5
C 2 H3 O 2 
0.1 M HC2H302
0.1 M NaC2H302
dihydrogen phosphate ion
H2P04
Ka = 6.2  10-8
HP042
0.1 M NaH2P04
0.1 M Na2HP04
bicarbonate ion
HC03
Ka = 4.8  10-11
C032
0.1 M Na HC03
0.1 M Na2C03
ammonium ion
NH4+
Ka = 5.6 x 10 -10
NH3
0.1 M NH4Cl
0.1 M NH3
Then calculate the ratio of volumes of conjugate base soluti on to acid solution:
Ka
[H 3O+]
=
V conjugate base soluton
V acid solution
Decide on the exact volume of buffer solution you will make (anywhere from 20 to 30 mL) and
calculate the volumes of acid solution and conjugate base solution you will use to prepare the buffer.
Show your calculations in part D on the report sheet (items 8 – 11 on page 8) and ask your
instructor to check them before you prepare your buffer solution.
Prepare the buffer solution and ask your instructor to verify your measured pH.
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6
REPORT SHEET
Experiment 16: Buffer Solutions
Name __________________________________________
last
first
Instructor’s Initials ___________
A.
AMMONIUM ACETATE BUFFER SOLUTION + STRONG ACID
1.
pH of 1.0 M NH4C2H3O2 solution
2.
Tube #1:
_______________
5 mL deionized water + 2 drops methyl orange indicator:
_________ drops of 1.0 M HCl solution needed to change pH of solution from 7 to 3.
3.
Tube #2:
5 mL 1.0 M NH4C2H3O2 (buffer) solution + 2 drops methyl orange indicator:
_________ mL of 1.0 M HCl solution needed to change pH of solution from 7 to 3
4.
Net ionic equation for the reaction of HCl and NH 4C2H3O2:
____________________________________________________________________________
B.
AMMONIUM ACETATE BUFFER SOLUTION + STRONG BASE
1.
pH of 1.0 M NH4C2H3O2 solution (from A-1, above)
2.
Tube #1:
_______________
5 mL deionized water + 2 drops Alizarin yellow R indiator:
_________ drops of 1.0 M NaOH solution needed to change pH of s olution from 7 to 11.
3.
Tube #2:
5 mL 1.0 M NH4C2H3O2 (buffer) solution + 2 drops Alizarin yellow R indicator:
_________ mL of 1.0 M NaOH solution needed to change pH of solution from 7 to 11
4.
Net ionic equation for the reaction of Na OH and NH4C2H3O2:
_____________________________________________________________________________
C.
SODIUM BICARBONATE BUFFER
1.
pH of 0.10 M NaHCO 3 solution
2.
Tube #1:
_______________
5.00 mL 0.10 M NaHCO 3 + 2 drops methyl orange indicator:
a.
_________ mL of 1.0 M HCl solution needed to change pH of solution to ~3
b.
Net ionic equation for reaction of HCl and NaHCO 3:
__________________________________________________________________________
3.
Tube #2:
5.00 mL 0.10 M NaHCO 3 + 2 drops Alizarin Yellow R indicator:
a.
_________ mL of 1.0 M NaOH solution needed to change pH of solution to ~ 11
b.
Net ionic equation for the reaction of NaOH and NaHCO3:
__________________________________________________________________________
7
D. PREPARATION OF A BUFFER SOLUTION:
1
Assigned pH:
___________________
2.
Calculated [H3O+]:
___________________
3.
Ka of Acid chosen:
___________________
4.
Ionization Reaction of Chosen Weak Acid
_________________________________________________ ____________________________
5.
Equilibrium Constant Expression:
6.
Formulas of Acid and Conjugate Base: _____________________and _____________________
7.
Calculation of Volume ratio of Conjugate Base to Acid in Buffer:
V conjugate base soluton
V acid solution
= __ Ka__
[H 3O+]
=
=
8.
Total volume of buffer solution to be made: _______________
9.
Calculation of Volume of Weak Acid Solution Needed to Prepare Buffer Solution:
1
10. Calculation of Volume of Conjugate Base Solution Needed to Prepare Buffer Solution:
(This will be a solution of a salt containin g the conjugate base of the weak acid.)
11. Buffer prepared by mixing ______________ mL of 0.1 M _______________________________
with ______________ mL of 0.1 M _______________________________________
Instructor’s approval __________
12. Measured pH of Buffer Solution _____________
Instructor’s approval __________
8
EXERCISES
1.
Write net-ionic equations for reactions that occur when each of the following is added to a buffer
solution composed of of Na 2HPO4 & NaH2PO4
a.
HBr solution added to the buffer solution:
_____________________________________________________________________________
b. b. KOH solution added to the buffer solution:
_____________________________________________________________________________
2. Write the formulas of the solute particles present in each of the following solutions. Classify each
particle as a Bronsted-Lowry acid and/or base. If more than one Bronsted -Lowry acid is present, list
the strongest one first. Decide on which of the solutions below would act as buffers (able to both
donate and accept protons).
Solution of
Solute Particles Present
in the Solution
Solute Particle that is
Brönsted base
(Proton Accepter)
Solute Particle that is
Brönsted Acid
(Proton Donor)
Buffer
(yes,
or no?)
NH4HSO 4
NH3 & NH4Cl
HCl & NaCl
NH4Cl & NaCl
H2CO3 & KHCO 3
HC2H3O2 & HCl
H2C2O4 & KHC 2O4
KHS & Na 2S
3. Calculate the pH of a solution prepared from equal volumes of equimolar solutions of H 2S and NaHS.
Ka1 for H 2S is 1.1 x10-7.
9
4.
Write the definition of a buffer.
5.
What effect does dilution (addition of water) have on the pH of a buffer solution? Will the pH
increase, decrease, or remain the same?
___________
6. Complete the table below. Write the formulas of solute particles present in each of the solutions. (In
#4,5, 7, 8, and 9, a double replacement reaction occurs between the compounds. Write the
balanced equation for the reaction and then, in the Solute Particles Present column, list the solute
particles remaining after the double replacement reaction has occurred.) in the solutions solutions.
Decide whether the solution would act as a buffer
Solute Particles
Present in the
Solution
1
Na2CO3
2
NaHSO 4
3
NaF
Solute Particle that is Solute Particle that Buffer
Brönsted base
is Brönsted Acid (yes, or
(Proton Accepter)
(Proton Donor)
no?)
Equal volumes of 0.10 M HCN and 0.05 M
NaOH. Equation for D.R. Reaction:
4
Equal volumes of 0.05 M H2S and 0.10
M NaOH. Equations for D.R. Reactions:
5
6
NaHC2O4
Equal volumes of 0.10 M NH3 and 0.10 M
HCl.
Equation for D.R. Reaction:
7
8
Equal volumes of 0.10 M NH3 and 0.05 M
HNO3.
Equation for D.R. Reaction:
Equal volumes of 0.10 M H2CO3 and 0.10
M KOH.
Equation for D.R. Reaction:
9
10