Hillerson 1
Natalie Hillerson
TA: Jeremy Lehner
21 January 2014
Buffers
I.
Introduction
The overarching purpose of this lab was to determine the buffering capacity of solutions
containing different phosphate species, specifically HPO42- and H2PO4-, and to compare
solutions prepared by dissolving aspirin and bufferin in water. The main concept in the lab
was the buffer: a solution containing a weak acid (or a weak base) and its conjugate base (or
conjugate acid). A good buffer will resist changes in pH upon the addition of a strong acid or
base. The quality of a buffer is determined by its buffering capacity, which is defined
quantitatively through the following equation:
∆[H + ]
𝛽=
∆pH
In this equation, ∆[H+] is the concentration of H+ ions added to the buffer solution, and ∆pH
is the subsequent change in pH. A good buffer will have a buffer capacity, as defined by this
equation, which is fairly large (less than 1, but high compared to other, less effective
buffers), meaning that the concentration of added hydrogen ions is smaller than but close to
the change in pH, indicating that the buffer successfully resisted pH changes while protons
were being added. A similar equation can be used for the addition of a strong base simply by
replacing ∆[H+] with ∆[OH-]. These equations were used for both the phosphate species and
the aspirin and bufferin comparisons.
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Buffers can be made with weak polyprotic acids, as was the case in this lab. Polyprotic
acids are amphoteric, meaning the ions of a polyprotic acid can behave as both an acid and a
base (accepting as well as donating protons) in successive deprotonations. The polyprotic
acid used in this lab was H2PO4- and HPO42-, both of which are represented by the following
chemical reactions:
HPO42- + H2O ⇋ OH- + H2PO4- Kb = 1.61 x 10-7
H2PO4- + H2O ⇋ H+ + HPO42- Ka = 6.23 x 10-8
For the first equation listed, HPO42- acts as a base because it will dissolve partially in water to
produce hydroxide ions and H2PO4-. This is why the Kb value was used as opposed to the Ka
value, and this Kb value was found by dividing Kw (1.0 x 10-14) by the listed Ka value of the
H2PO4- / HPO42- weak acid/conjugate base pair.
II.
Procedure
First, 100 mL solutions of both 0.1 M H2PO4- and 0.1 M HPO42- were prepared by
dissolving 1.3608 g of KH2PO4 and 1.742 g of K2HPO4 in 100 mL of water, respectively.
Then, 40 mL of the HPO42- was transferred into a separate beaker and 60 mL of DI water was
added, making solution 1. At this point, the pH probe was calibrated to samples of 4.00 pH
and 10.00 pH. The pH of solution 1 was then measured and recorded. Next, 20 mL of this
solution was again transferred into a separate beaker and .25 mL of .02 M HCl was added,
the pH then recorded. After disposing of this solution in the waste container, another 20 mL
sample of solution 1 was taken, but this time .25 mL of .1 M NaOH was added, the pH again
recorded. This 20 mL solution, as well as solution 1 was then discarded in the appropriate
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container. This same procedure was followed for solutions 2-5, according to the Figure 1 in
the Results section.
For the second part of the lab involving aspirin and bufferin, 0.2 g of both substances
were weighed and transferred to their own 125 mL beaker. The active ingredients of each
substance were recorded. Then, 25 mL of DI water was added to each beaker, and the
resulting solution was stirred. The solubility of each substance was also noted. After stirring,
the pH of the two solutions was recorded. To each solution, 0.25 mL of 0.5 M HCl was
added and the pH recorded. This entire process was then repeated, only adding 0.25 mL of
0.5 NaOH instead of HCl.
III.
Discussion
The phosphate-containing solution that had the largest buffer capacity overall was
solution 3, as shown in Figure 2, with the same mole amounts of H2PO4- and HPO42-.
Because this solution contained equal amounts of an acid and its conjugate base to begin
with, it more readily and easily reacted with the added protons and hydroxide ions and
essentially “neutralized” this addition to a good degree. As seen in the lab, solution 3 did
have the highest buffering capacity of any of the other solutions, with the exception of
solution 4—the addition of OH- resulted in a calculated high buffer capacity. This may be
due to human error in calculating the exact molarity of OH- ions to be added and actually
adding a different, lower concentration, which resulted in the skewed calculation of buffer
capacity.
The solution with the lowest buffer capacity was solution 1. In this solution, only HPO42ions are present (dissociation of this substance into H2PO4- is assumed to be negligible due to
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the fact that HPO42- is a very weak base). Since only a weak base exists, and not its conjugate
acid, the addition of H+ or OH- will change the pH more dramatically compared to solution 3,
the better buffer. Solution 5 is also presumed to be an ineffective buffer, but to a somewhat
lesser extent than solution 4. Although solution 5 was not tested in lab, subsequent
predictions and calculations indicate that it would also be a poor buffer for the same reasons
listed for solution 1—there is only H2PO4- present in solution.
Both the aspirin and bufferin solutions had very high buffering capacities—higher than
any of the phosphate-containing species. Aspirin was a better buffer than bufferin, if only
slightly. This is likely because even though the gram amount of both substances was the
same (0.2 g), the chemical makeup of these two compounds is different, and thus there were
different mole amounts in solution. Bufferin has additional inactive ingredients than aspirin,
so aspirin was shown to be the better buffer perhaps because the dissolved aspirin was more
concentrated in acid/conjugate base pairing, whereas the dissolved bufferin solution had
added inert species present in solution, making the acid/conjugate base pairing more dilute
and less effective as a buffer. Additionally, it is difficult to compute a theoretical buffering
capacity value for aspirin and bufferin, because both substances have complex chemical
makeups with many components that may affect their buffering capacities in ways that are
unknown and cannot be mathematically evaluated. Still, theoretically and experimentally,
aspirin has the higher buffer capacity than bufferin.
The experimentally calculated values of buffering capacity for the addition of H+ ions to
the phosphate-containing solutions compare quite well to the theoretical values of β, as seen
in Figure 3. For the majority of the solutions, the magnitudes of the values were the same,
and the numerical values were quite close as well. The only discrepancy evident was solution
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5. The theoretical value was quite small compared to the experimental value, and this might
be because the experimental value for solution 5 was an approximation, as this data was not
collected in the lab.
According to Figure 4, the pH range in which the buffering capacity is highest is around
7 and a bit higher, perhaps ranging from 7-8. This range of pH makes sense, as the pKa of the
H2PO4- / HPO42- weak acid/conjugate base pair is 7.21, and according to the HendersonHasselbalch equation (pH = pKa + log([A-]/[HA]), the pH equals the pKa when the ratio of
acid to conjugate base is equal to 1. A highly effective buffer forms when these two
concentrations are equal to 1, and as such, the pH of this buffer in question will be 7.21. For
the first deprotonation of H3PO4, the acid/conjugate base pairing is H3PO4 / H2PO4-, with the
pKa value being 2.12. Therefore, the most effective pH for a buffer containing H3PO4 and
H2PO4- would be 2.12. Likewise, for the pairing of HPO42- and PO43-, which has a pKa of
12.32, the most effective pH for this buffer would be 12.32.
IV.
Conclusion
This lab dealt entirely with buffers and their respective buffer capacities. This lab
demonstrated that the best buffers result from equal amounts of acid and conjugate base
present in solution before the addition of a strong acid or base (solution 3 is an accurate
demonstration of this statement). Likewise, those solutions with uneven concentrations of
acid and conjugate base at the start proved to be poor buffers with small buffer capacities.
Similarly, aspirin and bufferin both made highly effective buffers. Aspirin had a higher
buffer capacity than bufferin because its chemical makeup was less complex and therefore
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had more dissolved acid/conjugate base than bufferin, which had more dissolved species not
associated with buffers.
V.
Calculations
-
Calculation of experimental buffer capacity:
Solution 1: ∆[H+] = M1V1 = M2V2 = (.02)(.00025) = (x)(.02025)
∆pH = |initial pH – pH after addition of acid| = |9.06-8.83| = .23
∆[H+] / ∆pH = experimental buffer capacity, β = 1.07 x 10-3
Similar calculations were done for [OH-] and for solutions 2-5, DI water, aspirin, and
bufferin.
-
Calculation of theoretical buffer capacity:
β = ln(10) x ((Ka x [H+]) / (Ka + [H+])^2)) x C, where C = [HA] + [A-]
Solution 1: Ka = 6.23 x 10-8, [H+] = 10-8.83 (8.83 is the pH after addition of HCl), C = .04
M (dissociation is assumed to be negligible).
Using these values, β = 1.43 x 10-3
Similar calculations were done for solutions 2-5.