BCH 101- 1: Buffers
Buffers
Buffered solutions are quite important in chemical and biological systems. A buffer allows for the
maintenance of a fairly narrow range of pH even while another reaction is producing acids or
bases. Because a buffer is a mixture of a weak acid and its conjugate base, it can react with either
acid or base to remove the acid or base from solution. The functional pH range of a buffer is
described in the Henderson-Hasselbalch equation.
In order to solve problems with buffered solutions, it is essential to have an intuitive grasp of which
acid-base reactions are most likely to occur. In this section, we will learn how to determine which
reactions are most likely to occur in buffer problems. The following material is focused on problem
solving and will require close attention to details in problems and solutions presented.
Terms
Acid - A substance that has the potential to donate a proton or
accept an electron pair.
Acidic - Having a pH less than 7.
Base - A substance that can accept a proton, release OH-, or
donate an electron pair.
Basic - Having a pH greater than 7.
Buffer - A solution composed of an acid and its conjugate base
that serves to moderate the pH of the solution.
Conjugate Acid - A molecule that can be described as a base
that has gained one proton.
Conjugate Base - A molecule that can be described as an acid
that has lost one proton.
Dissociate - Separate into its ion constituents.
Indicator - A molecule whose conjugate acid or conjugate base
has a different color.
pH - A measure of the hydrogen ion concentration, it is equal to log [H+].
pK a - A measure of the strength of an acid, it is equal to – log K
a, where K a is the acid dissociation constant in water.
pK b - A measure of the strength of a base, it is equal to – log K b,
where K b is the base dissociation constant in water.
Strong Acid - An acid with a pK a less than zero. Strong acids
completely dissociate in water.
Strong Base - A base with a pK b less than zero. Strong bases
completely dissociate in water.
Titration - An experiment that neutralizes an unknown amount of
acid or base with a known volume and concentration of acid or
base to determine the amount of unknown acid or base.
Weak Acid - An acid with a pK a greater than zero. Weak acids
do not completely dissociate in water.
Weak Base - A base with a pK b greater than zero. Weak
bases do not completely dissociate in water. Buffered Solutions
How Buffers Work
As you have seen in calculating the pH of solutions, only a small amount of a strong acid is
necessary to drastically alter the pH. For certain experiments, however, it is desirable to keep a
fairly constant pH while acids or bases are added to the solution either by reaction or by the
experimenter. Buffers are designed to fill that role. Chemists use buffers routinely to moderate the
pH of a reaction. Biology finds manifold uses for buffers which range from controlling blood pH to
ensuring that urine does not reach painfully acidic levels.
A buffer is simply a mixture of a weak acid and its conjugate base or a weak base and its
conjugate acid. Buffers work by reacting with any added acid or base to control the pH. For
example, let's consider the action of a buffer composed of the weak base ammonia, NH3, and its
conjugate acid, NH4 +. When HCl is added to that buffer, the NH3 "soaks up" the acid's proton to
become NH4 +. Because that proton is locked up in the ammonium ion, it proton does not serve to
significantly increase the pH of the solution. When NaOH is added to the same buffer, the
ammonium ion donates a proton to the base to become ammonia and water. Here the buffer also
serves to neutralize the base.
As the above example shows, a buffer works by replacing a strong acid or base with a weak one.
The strong acid's proton is replaced by ammonium ion, a weak acid. The strong base OH- was
replaced by the weak base ammonia. These replacements of strong acids and bases for weaker
ones give buffers their extraordinary ability to moderate pH.
Calculating the pH of Buffered Solutions
Buffers must be chosen for the appropriate pH range that they are called on to control. The pH
range of a buffered solution is given by the Henderson- Hasselbalch equation. For the purpose of
the derivation, we will imagine a buffer composed of an acid, HA, and its conjugate base, A -. We
know that the acid dissociation constant pK a of the acid is given by this expression:
The equation can be rearranged as follows:
Taking the -log of this expression and rearranging the terms to make each one positive gives the
Henderson-Hasselbalch equation:
Note that the sample species HA and A- in the above Expression are generalized to the terms acid
and base, respectively. To use the equation, place the concentration of the acidic buffer species
where the equation says "acid" and place the concentration of the basic buffer species where the
equation calls for "base". It is essential that you use the pK a of the acidic species and not the pK b
of the basic species when working with basic buffers--many students forget this point when doing
buffer problems.
A buffer problem can be fairly simple to solve, provided you don't get confused by all the other
chemistry you know. For example, let's calculate the pH of a solution that is 0.5 M acetic acid and
0.5 sodium acetate both before and after enough SO3 gas is dissolved to make the solution 0.1 M
in sulfuric acid. Before the acid is added, we can use the Henderson-Hasselbalch equation to
calculate the pH.
This part of the problem does not require us to do the sort of equilibrium calculations that we must
use for Non-Buffered Solutions, but many students still try to do it the hard way. The hard way is a
correct way of doing the problem, but it may cost you valuable time on a test.
To calculate the pH after the acid is added, we assume that the acid reacts with the base in
solution and that the reaction has a 100% yield. Therefore, we say that 0.1 moles per liter of
acetate ion reacts with 0.1 moles per liter of sulfuric acid to give 0.1 moles per liter of acetic acid
and hydrogen sulfate. Here, we ignore the second dissociation of sulfuric acid because it is minor
in comparison to the first. So the final concentration of acetic acid is 0.6 M and acetate is 0.4M.
Plugging those values into the Henderson-Hasselbalch equation gives a pH of 4.57. Note that a
0.1 M solution of strong acid would give to a pH of 1 but the buffer gives a pH of 4.57 instead.
To probe the useful range of the buffer, let's calculate the pH of the solution resulting from the
same situation above but with different concentrations of the buffer. If the buffer is 1.0 M in both
acetate and acetic acid, then the pH of the resulting solution after the introduction of acid is 4.66.
However, if we make the solution only 0.11 M in acetic acid and acetate, then we calculate a pH of
3.45! Therefore, if you want a more effective buffer, make sure that the concentration of the
buffering agents is large in comparison to the added acid or base.
Problems and Solutions
Problem :
What is the ratio of base to acid when pH = pK a in a buffer? How about when pH = PK a + 1?
Solution: pH = pK a when the ratio of base to acid is 1 because log 1 = 0. When log (base/acid)
= 1, then the ratio of base to acid is 10:1.
Problem :
Explain why the pK a of a buffer should be as close as possible to the desired pH.
Solution: The pK a should be quite close to the desired pH so that the ratio of base to acid in
the Henderson-Hasselbalch equation will be close to 1. As the ratio of base to acid deviates from
1, the addition of acids and bases to the buffer will have a more profound effect on the pH.
Problem :
What is the pH of a buffered solution of 0.5 M ammonia and 0.5 M ammonium chloride when
enough hydrochloric acid is dissolved to make it 0.15 M HCl? The pK b of ammonia is 4.75.
Solution: The pK a of ammonium ion is 9.25 since pK a = 14 - pK b. 0.15 M H+ reacts with 0.15
M ammonia to form 0.15 M more ammonium. Substituting the values of 0.65 M ammonium ion
(acid) and 0.35 M remaining ammonia (base) into the Henderson-Hasselbalch equation gives a pH
of 8.98.
Terms