Review Experiment 3.
Answer Pre-Lab question posted in Pre-Lab Assignment folder (Experiment 3) on UM Learn.
The data sheet will become available for downloading on UM Learn at the beginning of the
lab session. To locate it, go to Content > Experiment 3 > Data Sheet Experiment #3.
Throughout the lab session you will be assigned to Break-out rooms to work on the questions.
You are encouraged to consult other students.
The data sheet is due at the end of this lab session (30 min after). You will receive an email
from Crowdmark at the beginning of the lab. Submit your data sheet following this link.
You will be marked on the answers to the questions provided. No formal report is required.
Your writing must be clear and legible. Writing that is too light to read or is illegible will
receive no credit.
To use a pH meter to investigate the buffer capacity for the addition of a strong acid and a
How do buffers resist changes in pH?
Chapter 15, Chapter 16 Sections 16.1-16.3 in Tro, Nivaldo J; Fridgen, Travis D.; Shaw, Lawton E.
Chemistry: A Molecular Approach 3nd Canadian ed.; Pearson Canada Inc.: Toronto, Canada, 2020.
Buffers are important for a wide range of disciplines, such as chemistry, biochemistry, biology,
agriculture, etc. In biological systems, for example, buffer solutions are often used to maintain a
particular pH required for the functioning of specific cellular processes. The buffering capacity in
the environment can determine whether an ecosystem will survive acid rain. The ability of buffers
to maintain the physiological pH allows the prevention of microbial growth. Also, they help to
increase shelf life and contribute to the efficacy of a drug. Contact lens solutions, skincare
products, baby wipes are everyday examples of buffer solutions.
In this experiment, we will prepare various buffers and compare their resistance to pH changes.
Buffers are chemical systems that resist changes to pH because they have a significant amount of a
weak acid and its conjugate weak base.
The pH of a buffer system is controlled by the equilibrium between the weak acid (HA), its
conjugate base (A−) and H3O+ concentrations:
HA (aq) + H2O (l)
A– (aq) + H3O+ (aq)
[A – ][H 3 O + ]
To find the pH of a buffer solution, we work an equilibrium problem using the ICE table in which
the initial concentrations include both the acid and its conjugate base.
Change in concentration
[HA]initial – x ≈ [HA]initial
[A–]initial + x ≈ [A–]initial
The “x is small” approximation is generally sufficient when [HA]initial and [A–]initial are within a
factor of 10 of each other, and both are at least 2 orders of magnitude greater than the
concentrations of [H3O+] and [OH–].
To simplify the finding of the buffer pH, we can use the Henderson-Hasselbach equation (2)
which allows a quick calculation of the pH from the initial concentrations of the buffer
components. This equation can only be used when the x is small approximation is valid.
[H 3 O + ] = K a
[A – ]
æ [HA] ö
[A – ]
– log [H 3 O + ] = – logçç K a – ÷÷ = – log K a + log
è [A ] ø
[A – ]
pH = pK a + log
Buffers can resist the pH change, because the acid (HA) and its conjugate base (A−) can readily
react with the strong base and strong acid, respectively.
When a strong acid is added to a buffer, it undergoes the following reaction:
H3O+ (aq) + A− (aq) ® HA (aq) + H2O (l)
Keep in mind that the addition of a strong acid is the same as adding H3O+ to the buffer. The
addition of H3O+ should lower the pH, but this additional amount will be used up through reaction
with the conjugate base, A−. Since A− reacts with the strong acid, there is less H3O+ in the solution
than we would expect and the pH change is very small.
The buffer system behaves similarly when a strong base, such as NaOH, is added. OH– reacts
with the acid component of the buffer (HA) and some of the acid is converted to the conjugate
base according the following reaction:
OH– (aq) + HA (aq) ® A– (aq) + H2O (l)
Although the addition of the base initially increases the [OH–] in solution, the resulting solution
has less OH– and the pH change is small.
By considering the amounts of a strong acid/base added, we can determine the new pH by either
using the Henderson-Hasselbach equation or by determining the excess amount of [H3O+] or [OH–]
remaining in solution.
To determine the effectiveness of a buffer, we use the buffer capacity. It is defined as the amount
of acid or base that can be added to a buffer without causing a large change in pH. The buffer
capacity increases with increasing absolute concentrations of the buffer components.
The buffer capacity (b) is defined as:
where DCa is the concentration of strong acid in the buffer solution, !
!”#$% “‘ ( %)*”+, (-./ (//$/
)”)(# 0″#1!$ “‘ %”#1).”+
and |DpH| is the change in pH, the absolute value of (final pH – initial pH).
Similarly, the buffer capacity with respect to strong bases can be determined using:
!”#$% “‘ ( %)*”+, 2(%$ (//$/
where, DCb is the concentration of strong base in the buffer solution !
)”)(# 0″#1!$ “‘ %”#1).”+
and DpH is the change in pH.
The greater the buffer capacity of a buffer solution, the more resistant it is to a pH change and the
more effective the buffer.
Buffer Choice and Preparation.
A buffer is most effective (most resistant to pH changes) when the concentrations of acid and
conjugate base are high and do not differ by more than a factor of 10.
Each buffer system works optimally over a specific pH range, and depending on the target pH,
different buffers can be used.
The pH range over which a buffer optimally resists changes is near the pKa of the acid. For example,
the pKa of acetic acid is 4.76, and so the optimal buffering ability occurs roughly around a pH of
4.76 ± 1.
You will be studying two different buffer systems in this lab. The first buffer is an acetic acid/acetate
buffer. This buffer tends to be effective in the pH range of 3.76 to 5.76. The chemical reaction for
this buffer is shown below:
CH3COOH (aq) + H2O (l) ⇌ CH3COO− (aq) + H3O+ (aq)
The second buffer is an ammonia/ammonium ion buffer, which tends to be effective in the pH range
of 8.24 to 10.24:
NH4+ (aq) + H2O (l) ⇌ NH3 (aq) + H3O+ (aq)
Ka = 5.8×10-10
Baby Wipes as a Buffer.
As mentioned previously, the ability of buffers to maintain the physiological pH has a lot of
benefits. For example, scientific studies show that baby’s skin has a naturally healthy pH between
4.5 and 6. But mess from dirty diapers can increase the pH making the skin more susceptible to
irritants resulting in diaper rashes. The key point to prevent irritation is maintaining a lower pH in
the diaper area. Baby wipes manufacturers claim that their products contain a specially designed
solution with a high buffer capacity that keeps the skin’s naturally healthy pH.
We will determine the buffering capacity of the buffer solutions extracted from two different brands
of baby wipes.
In the lab, pH is measured using a pH meter. The pH probe contains a reference solution with
known [H3O+], which is separated from the solution whose pH you want to measure by a glass
membrane. The pH meter measures the electrical potential difference between the solution being
measured and the reference solution within the pH probe. This electrical potential difference
arises from the difference in [H3O+] on either side of the glass membrane, as is the case in
concentration cells. A second electrode within the pH probe serves as a reference electrode that
has a precisely known potential.
® Calibration of pH meter
1. You will need three standard buffers: pH = 4.00 (red colour), pH = 7.00 (yellow colour), and
pH = 10.00 (blue colour).
2. Several brands of pH meters are provided in the lab. Usually, instructions for calibrating the
pH meter are provided on the meter itself. We will show you how to calibrate a pH meter
using the standard solutions.
® Preparation of Buffer Solutions
• Acetic acid/acetate buffer I:
10.0 mL 0.50 M acetic acid
0.41 – 0.42 g sodium acetate
90.0 mL deionized water
• Acetic acid/acetate buffer II:
10.0 mL 0.50 M acetic acid
0.11 – 0.12 g sodium acetate
90.0 mL deionized water
• Ammonia/ammonium buffer III:
10.0 mL 0.50 M aqueous ammonia
0.26 – 0.27 g ammonium chloride
90.0 mL deionized water
1. Use a 10-mL graduated cylinders to add the indicated amount of acetic acid or aqueous
ammonia to 250-mL beakers.
2. Weigh the indicated amounts of solid sodium acetate and ammonium chloride in a paper
weighing cup. Record the exact mass of the solids in your notebook to the nearest 0.001 g.
Add the solids to the beakers.
3. Use a 100-mL graduated cylinder to add 90.0 mL of deionized water to bring the final buffer
volume to 100.0 mL. Stir the solutions with a stir rod to thoroughly mix the solutions and
dissolve the solids.
4. You will be asked to calculate the number of moles of acetic acid, acetate, ammonia, and
ammonium in the buffer solutions. For this calculation, use the molarity of the acetic acid
and aqueous ammonia on the reagent bottle; use 82.03 g/mole for the formula mass of
sodium acetate; and 53.49 g/mole as the formula mass for ammonium chloride.
® Determining Buffer Capacity.
You will determine the buffering capacity of each of your solutions and compare this to the
buffering capacity of an unbuffered solution (deionized water).
The effect of Adding a Strong Acid (HCl).
Support a 50-mL buret on a ring stand with a buret clamp. Add 0.10 M HCl in a 250-mL
beaker. Fill the buret with 0.10 M HCl before setting it over the beaker containing buffer.
Measure 40.0 mL of the acetic acid/acetate buffer I in a 50-mL graduated cylinder and
transfer to a 100-mL beaker. Add a magnetic stir bar to the beaker.
Rinse the electrode with deionized water, dry it gently with a KimWipe and place it in the
buffer solution. Do not allow the stir bar to hit the electrode. Measure and record the
initial pH of your buffer and the initial buret reading.
While stirring the solution continuously, slowly add HCl dropwise from the buret.
Continue adding the HCl until the pH changes by at least 1 pH unit. Record the final pH
and buret reading in your data sheet.
Repeat this process for acetic acid/acetate buffer II, ammonia/ ammonium buffer III and for
The effect of Adding a Strong Base (NaOH).
Add 0.10 M NaOH in a 250-mL beaker. Rinse and fill a second 50-mL buret with 0.10 M NaOH.
Using the same protocol, add 0.10 M NaOH to 40 mL samples of each buffer and 40 mL of
deionized water. Remember to record the initial pH.
® Buffer Capacity of Baby Wipes.
We will prepare additional buffers using baby wipes and determine the buffering capacity of the
resultant solution1. There will be two brands of baby wipes to study. Record the brand/type of
baby wipe in the data sheet.
1. Place eight baby wipes in 600 mL beaker. Use a 100-mL graduated cylinder to measure
90.0 mL of deionized water. Add water to the beaker of wipes and allow the wipes to
become saturated with water. Allow the wipes to sit for 5 minutes.
2. Gently stir the wipes with a spatula for one minute, then one at a time, remove each baby
wipe from the beaker and wring dry by twisting and squeezing the baby wipe. Collect the
resulting solution in the same 600 mL beaker.
3. Measure 40.0 mL of the baby wipes solution in a 50-mL graduated cylinder. Transfer the
solution to a 100-mL beaker.
4. Use the same procedure as you did previously to measure the effect of adding HCl and
NaOH to the acetate and ammonium buffer, add acid and base dropwise and determine the
buffer capacity of the solution from the baby wipes. Please note, the concentration of the
buffer in the baby wipes is very low! You will need a very small volume of HCl and NaOH
to effect a large change in the pH. Add HCl and NaOH drop by drop and allow the pH to
stabilize between additions. Record your results.
1. J. Chem. Educ. 2018, 95, 1816-1820.
2. Tro, Nivaldo J; Fridgen, Travis D.; Shaw, Lawton E. Chemistry: A Molecular Approach
3nd Canadian ed.; Pearson Canada Inc.: Toronto, Canada, 2020.
While studying for an exam, a student starts breathing rapidly and feels lightheaded. A
classmate gets the student to breathe into a bag. He explains that the condition is called alkalosis
and causes the blood pH to get too high. Breathing into the paper bag (and breathing from the
bag) helps lower the blood pH.
Explain this fact.
® Chemistry and Medicine Box. Buffer Effectiveness in Human Blood (section 16.3, p.
705) Tro, Nivaldo J; Fridgen, Travis D.; Shaw, Lawton E. Chemistry: A Molecular
Approach 3nd Canadian ed.; Pearson Canada Inc.: Toronto, Canada, 2020.
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