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| Droppers Glassware Graduated Cylinders pH test paper Spot Plates Volumetric Flasks |
| A word about notation: "pKa" and "pKa" are sometimes used interchangeably in lab manuals. The more correct notation is pKa, because "K sub a" is the acid dissociation constant. "Ka" implies "K" times "a", which is not the case here. The Henderson - Hasselbalch equation is sometimes given with just "pK". The subscript "a" is assumed. |
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CAUTION: If you choose to attempt any of the experiments or procedures described on this site, you do so entirely at your own risk.
Experiment: Making a Buffer Solution
By Christian ThorstenIntroduction: This is technically an inorganic chemistry experiment, but acid-base buffers and buffering are crucial in the study of biochemistry. Many students don't learn to prepare buffers until they take a biochem lab course. Unfortunately, in college the Henderson-Hasselbalch equation often becomes just one more thing to memorize for an exam, while the true usefulness of buffers is realized only in subsequent lab work or graduate studies.
The deceptively simple equation has profound implications
for chemical and biological systems:
In this experiment, the weak acid "HA" will be orthoboric acid (H3BO3), while the salt "A" of this weak acid will be borax, Na2B4O7.10H2O. In stoichiometric terms, borax is not simply the sodium salt of H3BO3, but here it is treated as such for simplicity. "HA" is the protonated form of an acid, while "A-" is the ionized form; the H+ or "proton" is therefore detached. Strictly speaking, any given "A-" that's floating around in solution could have originated from the ionization of the acid or from the ionization of the buffer salt. Most commonly, the acid is paired with a buffer salt that gives the identical anion species (e.g., acetic acid and sodium acetate both give CH3COO- anions in solution).
Incidentally, H3BO3 isn't even a Bronsted acid (proton donor) in solution... it's a Lewis acid (electron pair acceptor); however, the boric acid-borax buffer still works, because boric acid can be thought of as a "virtual" proton donor. It abstracts an OH- from water and leaves an H+.
Supplies: Boric acid (orthoboric acid); Sodium tetraborate decahydrate (borax); Distilled-deionized water; Volumetric flask (250 mL); Laboratory balance; 0.01-molar (10 mM) Hydrochloric acid solution (provided by your lab instructor) ; Droppers; Wide-range pH paper or Wide-range pH indicator kit; Safety Goggles.
Safety: Borax/boric acid buffer itself is quite safe on the relative scale; just don't ingest it. Most common buffers have very low toxicity, but choice of buffer salts can impact this greatly (cacodylate comes to mind as one of the more toxic varieties).
0.01 molar HCl, used here to test the buffer's
effectiveness, is in the range of stomach acid
(0.01 M HCl = pH 2, with stomach acid being anywhere from 1 to
3). Don't get it in the eyes, and avoid contact with
skin or clothing. Stomach acid is
nowhere near
as concentrated as 6M or 12M HCl, but it can still cause burns;
that is why acid reflux disease is so troublesome. Keep those safety goggles on your eyes. The diluted
acid
should be prepared by your lab instructor so you won't have to
handle the concentrated HCl. For maximum safety, wear a lab
coat and preferably a rubberized lab apron. There is no point in
risking holes in one's clothes; these sometimes don't appear
until the clothing is washed.Procedure:
I. Preparing the buffer:
Boric acid: mol. weight 61.83 g.
Sodium tetraborate 10-hydrate: mol. weight 381.37 g.
Weigh out 0.77 gram of boric acid and 4.76 grams of borax; this corresponds to about 0.0125 mole of each. Add these to a 250 mL volumetric flask and add a small amount of de-gassed, de-ionized water, just enough to dissolve the solids. Swirl until the solids dissolve, then add enough distilled water to make the total volume up to 250 mL. The boric acid may take a while to dissolve, especially if it has formed lumps.
Based on the amounts used, the final concentration will be 0.05 M.
Pipette a
small amount of this solution and add 5 drops of it to a clean well in
the spot plate. Add 1 drop of
wide range
pH indicator solution to this (or put one drop on
a
piece of wide range pH test
paper) and immediately note the color. Check the color against the
chart to
determine approximate pH. It should be around 9.0 to 9.5
1 II. Testing the buffer:
Measure
out 50. mL of the
buffer solution and put it into a beaker. Add 2 drops of
0.01-molar HCl. Stir with a glass rod,
taking
care not to spill any. Then, take five drops of the resulting solution
and
place in a clean well on the spot plate. Add 1 drop of wide range pH
indicator
and note the color. What is the approximate pH, according
to the
chart? If a calibrated pH meter is available, use it to test the
exact pH before and after adding the 0.01 M HCl.If you were to add the
same amount of 0.01-molar HCl to an unbuffered solution of pH 9.2 2 ,
what would you expect the new pH to be? 3 What difference, if any,
is there
between actual and
expected results?If, hypothetically, the experiment had been tried using
0.1-molar HCl instead of 0.01-molar HCl, would the buffer solution
still
"absorb" it without appreciably changing in pH? Why or why not? What
would happen if one were to use 1.0-molar HCl? How much HCl would be required to overwhelm the buffering
capacity?Discussion: As this simple experiment demonstrates, acid-base buffer solutions tend to resist changes in pH. This is essential to living organisms. Simple processes such as eating, drinking, or taking medicine would throw the body into fatal imbalance if it weren't for our natural buffering systems.
Ideally,
the pH at which
a solution is intended to buffer should be close to the pKa of the
buffering
species. At room temperature, boric acid has a pKa near 9.2; therefore,
a
borate-boric acid buffer solution should be expected to buffer most
effectively
at a pH between 8 and 10. In other words, a system designed to buffer
at pH 9 will maintain a pH at or close to 9, even if some acid or
alkali
is added.It is entirely possible, indeed common, to prepare buffers
that have the same pH but different molarity. While ionic
strength does affect matters, the buffering pH of a system depends much
more on the ratio of weak acid
to buffer salt than it does on the overall concentration of dissolved
compounds. (Buffering capacity,
on the other hand, depends greatly on concentration). The
Henderson-Hasselbalch equation helps us determine
what that ratio should be for any given system. The "pK" in the
equation is simply the pKa of the weak acid on which the
buffer is based. In our example, it's boric acid (H3BO3).
If we wanted a buffer with pH identical to the pK of the weak acid (a
desirable situation, as mentioned previously), we would want log ([A-]/[HA])
to equal zero, meaning the ratio of [A-] to [HA] would be 1.
(recall that 10 to the 0 power equals 1; conversely, log of 1
equals 0). If we treat the salt of a weak acid as something that
ionizes completely (as salts do), and the weak acid itself as something
that ionizes to only a negligible extent (not quite true, but adequate
for our calculations), we can therefore use a 1:1 mole ratio of weak
acid : salt of the weak acid.When testing the buffer, there is a good reason that
vinegar (CH3COOH) isn't used as the acid, nor bicarbonate
(HCO3-) or carbonate (CO3--)
as the base. Acetic acid and bicarbonate ionize only weakly,
adding some buffering effect of their own; thus, they would
complicate our simple experiment too much. Moreover, carbonate
and bicarbonate lose CO2 on exposure to acids. Notes:
1 this should be the case if there is no carbon dioxide dissolved in the water. Aqueous CO2 creates carbonic acid and lowers the pH. There's another problem: ionic strength of the buffer solution also has a great effect on pH. Theoretical buffer pH and real buffer pH are usually two different numbers. For best accuracy, it's recommended to use a pH meter. While monitoring the solution's pH, small amounts of either HCl or NaOH in dilute solution are added until the pH reaches the desired value.
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2 We'll assume the unbuffered solution is made pH 9.2 by preparing approx. 16 micromolar NaOH in water. Only a very strong base such as this can be thought of as having no buffer capacity. Weak bases such as Na2CO3 inevitably give the solution some buffering capacity.
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3 If we were using pure water (pH 7), the calculation of "expected" pH after adding acid or base would be more straightforward. When adding acid to a basic solution (or the other way around) and trying to calculate the final pH, one must of course take into account the amount of H+ that will combine with OH- to form water.
To the solution we added 2 drops, which translates to about 0.1 mL. In the case of 0.01M HCl that will be 1 micromole of H+ ions added to the solution. If the solution is pH 9.2 to begin with, we must calculate how many moles of OH- ions there are to react with any added H+. Subtraction then tells us how many moles of H+ this will leave. This will allow us to calculate the final pH of the unbuffered solution.
At pH 9.2, [OH-] is antilog{-(14 - 9.2)} = antilog(-4.8) = 1.6 x 10-5 moles per liter. 50 mL of this solution translates to 8.0 x 10-7 moles of OH-. If 1 micromole of H+ was added, that is 1 x 10-6 moles H+, minus 8.0 x 10-7moles that react with OH- (or 0.8 x 10-6 moles, to make calculation easier). That leaves 0.2 x 10-6 moles H+ (which equals 2x10-7 moles H+). This many moles dissolved in 50 mL of solution gives a [H+] of 4 x 10-6 M. The expected pH would be about 5.4 with no buffering.
According to a preliminary test run with wide-range pH paper, the buffered solution was still around pH 9 after addition of the same amount of HCl. That is a gigantic difference in hydrogen ion concentration: pH 5.4 versus pH 9.
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