Buffers and Buffering Capacity
Learning Goals
1. Investigate the buffering ability of a solution containing both a weak acid
and its conjugate base.
2. Prepare several buffered solutions and test their buffering capacity.
Abstract
The solutions to be investigated in this experiment are buffers. A buffer
is composed of a weak acid or base and the salt of the weak acid or
base, called the conjugate. For example, acetic acid is a weak acid and
in a dilute solution a small number of the acetic acid molecules
dissociate to produce acetate ions and protons, H+. The acetate ions
are the conjugate base of the acetic acid. Buffer solutions are made up
of an acid-base pair composed of either a weak acid and its conjugate
base or a weak base and its conjugate acid.
For example the following are acid-base pairs from which a wide
variety of Buffers can be made:
Table 1. Examples of Acid-Base pairs
Acid
Acetic Acid
CH3COOH
Monobasic
phosphate ion
H2PO4−
Citric acid
C3H5O(COOH)3
Ammonium ion
NH4+
Base
De-protonation reaction
Acetate ion
CH3COOH ⇔ CH3COO− + H+
−
CH3COO
Dibasic phosphate
H2PO4− ⇔ HPO42− + H+
ion
HPO42−
Monobasic citrate ion C3H5O(COOH)3 ⇔ C3H5O(COOH)2COO− + H+
C3H5O(COOH)2COO−
Ammonia
NH3
NH4+ ⇔ NH3 + H+
To be effective as a buffer, the solution must be composed of the correct acid-base
pair and their concentrations must be in the exact ratio that will give the desired pH
value. The prepared buffer solution also needs to have a sufficient concentration of
both the acid and the base so that the solution will have the strength to neutralize a
given amount of a strong acid or base. This strength of the buffer solution is called its
Buffering Capacity.
So a Buffer solution has the following three requirements:
1. It is made up of the correct acid-base pair to get close to the pH value needed.
2. It has the acid and base concentrations in the exact ratio for the exact pH value.
3. It has sufficient concentration of the acid-base pair to produce the required Buffering
Capacity.
A pH meter is a valuable tool that can be used to test the pH of prepared solutions and
gain a better understanding of buffers and their buffering abilities. Using the
understanding of the chemistry of buffered systems, buffer solutions can be prepared
that have a specific pH value and buffering capacity. A Buffered system with pH < 7
consists of a weak acid and a salt with the same anion as the weak acid. For pH > 7,
the buffer is made up of a weak base and the salt of that weak base.
For example: A solution containing a relatively large concentration of a weak acid, e.g.
Formic acid (HOOCH), and the corresponding salt of that weak acid, Sodium formate
(NaOOCH), is a buffer. As the weak acid and its conjugate base are dissolved into the
solution the following reactions occur:
HOOCH(aq) ⇔ H+(aq) + OOCH-(aq)
NaOOCH(aq)
⇒
Na+(aq) + OOCH- (aq)
(de-protonation)
(dissolution)
The addition of the common anions from the acid and its salt establishes an equilibrium
that has a high concentration of formic acid and its conjugate base, the formate ion, as
shown:
HOOCH(aq) ⇔ H+(aq) + OOCH-(aq)
This is the acid-base pair that makes up the buffer system. If strong acids are added to
the buffered solution, the strong acid will react and be neutralized by the conjugate base
present in the buffer. This drives the equilibrium of the above to the right, to reform
some of the conjugate base ions. If strong bases are added to the buffered solution, the
strong base will react and be neutralized by the weak acid present in the buffer. This
drives the equilibrium of the above reaction to the left, to reform some of the acid form of
the buffer pair. Therefore, the addition of acid or base to this buffered solution will not
cause a large change in pH, but will cause a change in the concentrations of both the
weak acid and the conjugate base.
Your text (Moore, Stanitski, Jurs) has an excellent explanation of Buffer calculations. See
Chapter 17,Section 1, Pages 805-816.
Prelab Assignment
In your lab notebook, prepare the following information:
1. A brief (2-3 sentence) introduction to the lab.
2. A table of safety information including the chemicals used in the lab and their
safety handling precautions. This information can be obtained from the
Material Safety Data Sheets (MSDS).
3. Pick a pH (9.0, 7.0, or 3.0) and using your knowledge of equilibrium, pH, Ka
and Kb values, choose a buffer pair from the acids, bases and salts provided in
Table 2 (below). Determine the correct molar ratio of acid or base to salt
needed for the selected pH. Then determine the mass (for solids) or volume
(for liquids) of the chemicals needed to make 100 mL of the buffer with a
Buffering Capacity so that 50 mL of the buffer is capable of neutralizing at
least 2 mL of either 1 M HCl or 1 M NaOH without letting the pH change by
more that 1 pH unit. See your text pages 815-816. You may also need to look
up the solubility of each chemical. See the MSDS forms for this information.
Here is an example of a solved buffer calculation and
Here is the same problem solved using the Henderson-Hasselbalch equation.
You will not be allowed to work in the lab without Pre-lab Calculation #3.
Table 2. Available Acid-Base pairs for Preparation of buffer solutions in this experiment
Chemical
Type
Formula
Molecular
Weight
Ka
pKa
Kb
pKb
4.75
−
−
−
−
−
-8
7.21
−
−
-13
12.32
−
−
-4
3.128
−
−
−
−
−
9.24
1.7 x 10
−
−
Acetic acid 1.0M
solution
Weak acid
HC2H3O2
60.05
1.8 x 10
Sodium acetate,
anhydrous
Conjugate
base of
above acid
NaC2H3O2
82.03
−
Sodium phosphate
monobasic
monohydrate
Weak acid
NaH2PO4 ŸH20
137.99
6.2 x 10
Sodium phosphate,
dibasic heptahydrate
Conjugate
base of
above acid
Na2HPO4 Ÿ7H20
268.07
4.7 x 10
Citric acid
monohydrate
Weak acid
C6H8O7 ŸH20
210.14
7.5 x 10
Sodium citrate
dihydrate
Conjugate
base of
above acid
Na3C6H5O7 Ÿ2H20
294.11
−
Ammonia 2.0 M
solution
Weak base
NH4OH
17.3
Ammonium chloride
Conjugate
acid of above
base
NH4Cl
53.5
5.9 x 10
−
-5
-10
-5
4.76
−
Chemicals and supplies
Chemicals
Liquids:
Deionized Water
1.0 M Acetic acid (HC2H3O2)
1.0 M Hydrochloric acid (HCl)
1.0 M Sodium Hydroxide (NaOH)
2.0 M Ammonia solution (NH4OH)
Solids:
Sodium acetate, anhydrous (NaC2H3O2)
Citric acid, monohydrate (C6H8O7 Ÿ H2O)
Sodium phosphate, monobasic monohydrate
(NaH2PO4 Ÿ H2O)
Sodium phosphate, dibasic heptahydrate
(Na2HPO4 Ÿ 7H2O)
Sodium citrate, dihydrate (Na3C6H5O7 Ÿ 2H2O)
Ammonium chloride (NH4Cl)
Glassware and Supplies
pH meter
1mL plastic syringes
100mL volumetric flask
magnetic stirrer and stir bar
Procedure:
Part A. Preparing an Acetic acid/Acetate Buffer
1. Weigh 0.41g of solid Sodium acetate, anhydrous into a beaker. Add 5.0 mL of 1M
2.
3.
4.
5.
6.
7.
8.
Acetic acid and about 50mL of deionized water to dissolve the solid. Quantitatively
transfer the solution to the 100mL volumetric flask and fill to the mark with more
deionized water. Calculate the concentrations of Acetic acid and Acetate ion.
Divide this solution equally between two beakers labeled B1 and B2.
Prepare two other beakers with 50 mL of deionized water in each. Label these
beakers D1 and D2.
Using a ring stand and clamps, assemble the magnetic stirrer and pH meter, so that
pH can be measured while stirring without the stir bar hitting the electrode.
Place the pH electrode in the solution labeled D1. Measure and record the initial pH.
Using a 1mL plastic syringe, add 0.2 mL of 1.0 M HCl, while stirring, and record the
pH. Repeat the addition of 0.2 mL more of HCL, while stirring, and record the pH
after each addition, until the pH changes dramatically (± 2 pH units).
Repeat Step 5 using the beaker labeled D2, only this time use the 1.0 M NaOH
instead of the HCL, recording the pH after each addition of NaOH.
Repeat Step 5 using the beaker labeled B1 (the buffered solution) and adding HCl.
Repeat Step 5 again using the beaker labeled B2 and adding NaOH.
Part B. Preparing a Buffer of Specific pH and Buffering Capacity
1. Prepare 100 mL of your buffer solution, pH 9.0, pH 7.0 or pH 3.0, based on your
Pre-Lab calculation.
2. Test the buffering capacity of your prepared buffer: evenly divide the 100 mL of your
buffer solution into two beakers, and add strong acid to one and strong base to the
other, recording the pH of the solution with each addition of acid/base (like steps 5
and 6 in Part A).
Rubric for Lab Report: Buffers and Buffering Capacity Claim(s):
Statement(s), derived
from evidence, using
scientific reasoning.
(15 points total)
Summary of claims:
Clearly state the major conclusions or claims for this experiment:
What buffering capacity was achieved with the acetic acid/acetate
buffer? With the other buffering system selected?
Evidence:
Scientific data that
supports the claim.
The data needs to be
appropriate and
sufficient to support
the claim. Submission
of the Analog to
Target and Designing
Experiments
Worksheets are
required.
(30 points total)
Introduction:
Provide background information to put the experiment in context:
What is a buffer and what is its purpose? What is meant by “buffering
capacity”?
Reasoning:
Scientific explanations
that use evidence and
appropriate chemistry
concepts to construct
claims.
(30 points total)
Discussion:
Explain how the evidence you presented supports your claim.
Show calculation of moles of acid/base added to the acetate buffer
system. Show calculations for how to prepare your specific buffer and
its experimentally determined buffer capacity. Did the buffer have
sufficient buffering capacity to add 2mL of strong acid/base and keep
the change in pH ≤ 1 unit? Explain any noted errors that occurred in
lab and how that affected your results.
Procedure:
In a narrative (not numbered directions) summarize the major steps of
the experiment. You may reference parts of this procedure by
providing the title and web address.
Results:
Organize and present all the data you collected. Include graphs of pH
vs. volume of acid/base added to D1, D2, B1 & B2. Include graphs of
pH vs. volume acid/base added to your specific buffer.
Copyright (c) 2014, the ICN Team.