Buffers: Applications in Chemical
Equilibrium
Minneapolis Community and Technical College
C1152: Principles of Chemistry 2
v.10.13
Introduction
A buffer is a mixture of a weak acid and its conjugate base, or a weak
base and its conjugate acid. The buffer’s function is to neutralize
additional acids (H3O+ ions) or bases (OH– ions) thus keeping the pH
of the solution approximately constant.
Of course, no buffer is perfect and slight changes in pH will be observed as the buffer operates within its buffering range
(figure above). As an acid is added to the buffering system, the buffer’s own base is consumed and the pH drops slightly.
Similarly, when strong base is added to the buffer, the buffer’s own acid is consumed and the pH rises slightly. Eventually,
either the buffer’s acid or base component is used up and the pH dramatically changes. At this point the buffer is no longer
operating within its “buffering range” and we say the buffer is exhausted. A buffer will have greater capacity if the
concentrations of the conjugate acid and base are high. This is because there are more buffering moles present to react with
additional strong acid or base. The buffer capacity is also greatest when the concentrations of the conjugate acid and base are
equal. In this case pHbuffer = pKa.and the buffer is positioned in the middle of its buffering range.
The pH range in which a buffer solution is effective is generally considered to be ±1 of the pKa value. In this experiment, we’ll
assume the buffer’s capacity to be determined by the same range of pH values.
In many chemical and biological systems buffers are important. Blood plasma, a natural example in humans, is a bicarbonate
buffer that keeps the pH of blood between 7.2 and 7.6. pH’s outside of this range are fatal.
A buffer is an equilibrium system. Consider the weak acid (nitrous acid, HNO2) equilibrium below:
HNO2 (aq) + H2O(l) 
H3O+ (aq) + NO2 – (aq)
For the forward reaction, HNO2 is acting as a Lowry Bronsted acid (proton donor) and the H2O molecule acts as the base
(proton receiver). For the reverse reaction the H3O+ ion acts as an acid and the NO2- acts as the base. To be a buffering
system, the equilibrium must possess significant concentrations of both the weak acid (HNO2) and the conjugate base
(NO2-).
The equilibrium constant for this reaction is written as:
where the bracketed quantities represent equilibrium concentrations.
To prepare a buffer system with nitrous acid, the conjugate base (NO2-) is added as solid
NaNO2. The sodium ions are spectator ions. The resulting buffering mixture contains
HNO2 molecules capable of neutralizing hydroxide ions and nitrite ions (NO2-) that
neutralize added acids. A variation of the equilibrium expression above, called the
Henderson-Hasselbalch equation (right), is the useful in preparing a buffer solutions.
Because the HH equation utilizes the x~0 approximation, it can only be used for relatively
dilute solutions. For our nitrous acid/sodium nitrate buffer example, the HendersonHasselbalch equation is shown at right where Ka = 4.60 x 10-4 for HNO2.
In this experiment, you will use the Henderson-Hasselbalch equation to determine the amount of solid sodium acetate
(NaCH3COO) you will add to 100mL of 0.100 M acetic acid (CH3COOH) to obtain a specific pH. You will then prepare the
buffer and titrate it with NaOH and HCl solutions to determine the buffer’s limitations. You will also titrate one of two
NH3/NH4+ buffers and then compare its behavior to the CH3COOH/CH3COO- buffering system.
PRE-LAB EXERCISE
Use the Henderson-Hasselbalch equation to calculate the mass of solid sodium acetate (NaCH3COO) required to mix with
100.0 mL of 0.100 M acetic acid (CH3COOH) to prepare buffers with pH = 4.9 and pH = 3.8. The Ka of acetic acid is
1.76 10–5. and the molar mass of sodium acetate is 82.0339 g/mol.
MATERIALS
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Vernier computer interface
magnetic stirrer and 2 cm stirring bar
Vernier pH Sensor
articulated arm (for pH probe)
pH = 2 and pH = 12 buffer solutions
in vials for probe calibration
0.500 M sodium hydroxide solution
0.500 M hydrochloric acid solution
0.100 M acetic acid solution
solid sodium acetate, NaC2H3O2
NH4+/NH3 Buffer solutions A and B
2 – 50 mL beakers
3 - 150 mL beaker s
2 - 25 mL graduated cylinders
1 - 100 mL graduated cylinder
two 25 mL burets and buret clamps
distilled water
Top loading balance
Water soluble marker to label burets
Kim Wipes
(HCl and NaOH solution transfer to burets)
(one for buffer prep and the other two for titration)
(buffer & distilled water measurements)
(0.100 M acetic acid measurement)
(HCl and NaOH dispensing)
PROCEDURE
Part I pH Probe Calibration
Calibrate your pH probe using pH = 2 and pH = 12
buffer solutions provided in capped vials. Remember
to keep the pH probe immersed in clean distilled water
when not being used.
Part II Buffer preparation and adjustment
You will be assigned a buffer pH when you come to
class. At that time you and your partner will be asked
to calculate the mass of solid NaCH3COO (82.034
g/mol) you will need to combine with 100.0 mL of the
0.100 M acetic acid for your assigned pH.
HCl Titration
Adjust each buret to read exactly zero.
Use a graduated cylinder to measure out 10.0 mL
of your buffer solution into a 150 mL beaker and
add 15.0 mL of distilled water. Place the beaker on
a magnetic stirrer and adjust the positions of the
buret and pH probe and magnetic stirrer.
Record the initial buret reading (zero) and
corresponding pH. Now carefully add 0.500 M
HCl solution to the buffer in 0.2 mL increments.
When you have completed your calculation show it to
the lab instructor who will check it for correctness.
After each addition, monitor the pH for 5 seconds
before recording the buret reading in your
notebook.
Next, use a top loading balance to weigh out the
sodium acetate in a tared 150 mL beaker. Dissolve the
solid in 100.0 mL of the 0.100 M acetic acid solution.
Continue adding HCl incrementally and recording
the pH until you have at least 8 points at very low
pH.
Measure the pH of the buffer solution with your
calibrated pH probe and record the pH in your
notebook.
Dispose of the used buffer solutions in the sink.
Part III: CH3COOH/NaCH3COO
Buffer titrations
Set up two burets using a single buret clamp. Rinse
and fill one buret with 0.50 M NaOH solution and the
other with 0.50 M HCl solution.
Adjust each buret to read exactly zero.
NaOH titration
Use a graduated cylinder to measure out 10.0 mL of
your buffer solution into a 150 mL beaker and add 15.0
mL of distilled water. Place the beaker on a magnetic
stirrer and adjust the positions of the buret and pH
probe and magnetic stirrer.
Record the initial buret reading (zero) and
corresponding pH. Now carefully add 0.500 M NaOH
solution to the buffer in 0.2 mL increments.
After each addition, monitor the pH for 5 seconds
before recording the buret reading in your notebook.
Continue adding NaOH incrementally and recording
the pH until you have at least 8 points at high pH.
Part IV: NH4+/NH3
Buffer titrations
Your instructor will assign your team one of two
different NH4+/NH3 buffer solutions (A or B) to
analyze.
Repeat the HCl and NaOH titration procedures
above for the NH4+/NH3 buffering system you’ve
been assigned. Don’t forget to dilute the 10 mL of
buffer with 15 mL of distilled water before
beginning.
Team Report
Page 1: Graphs
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Upper right corner: Your names, lab section number & Date of exp.
CH3COOH/NaCH3COO buffer titration graph (see figure at right).
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Y axis: pH
X axis: leftward mL of HCl (acid) and rightward mL of NaOH (base)
Excel: R Click on graph, Select Data, Add
Use negative HCl volumes to flip the HCl data.
Labeled pKa point in the exact center of the buffering range.
Graph title and axis labels
Horizontal +/- 1 pH lines (see description below)
Vertical buffer capacity lines (see description below)
NH4+/NH3 Buffer titration graph (see figure at right).
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Clearly indicate which buffer you were assigned (A or B) in the graph title.
Y axis: pH
X axis: leftward mL of HCl (acid) and rightward mL of NaOH (base)
Excel: R Click on graph, Select Data, Add
Use negative HCl volumes to flip the HCl data
Labeled pKa point in the exact center of the buffering range.
Graph title and axis labels
Horizontal +/- 1 pH lines (see description below)
Vertical buffer capacity lines (see description below)
A buffer’s optimum operational pH is at it’s pKa point. However, we consider
the buffering range to be +/- 1 pH unit on either side of the pKa point. Use your
pKa point to determine each buffer’s optimum pH range and list these pH ranges
on each graph’s upper left-hand corner. Now draw 2 horizontal lines on each
graph at locations +1 and -1 pH unit from their respective pKa points. (see figure
at right)
Buffer capacity has a rather loose definition, yet it is an important property of
buffers. A commonly seen definition of buffer capacity is: “The amount of H+ or
OH– that can be neutralized before the pH changes to a significant degree.”
Use your graph to determine each buffer’s capacity in total mL of HCl and
NaOH. To do this, draw vertical lines from where the horizontal pH limits
intersect the titration curve (figure at right). Determine the total mL between the
two limits and report this as the “Buffering Capacity” in the upper right hand corner of each graph.
Page 2: Questions
1. Which buffer has the greatest capacity? Why?
2. A buffer is constructed by mixing 15.46 g of sodium acetate, NaC2H3O2, with 100.0 mL of 1.50 M acetic acid.
The Ka of acetic acid is 1.76 10–5.
a. What is the initial pH of the buffer? (Show all work for credit)
b. What volume of 1.0 M NaOH would be required to reach the buffer’s equivalence point?
(Show all work for credit)
c. What is the pH of the buffer at the equivalence point? (Show all work for credit)
d. Calculate the pH of buffer C after a total addition of 175.0 mL 1.00 M NaOH solution. (Show all work for credit)