Content Reviewers:Rishi Desai, MD, MPH
Buffer solutions contain relatively high concentrations of a weak acid and its conjugate base.
When small amounts of acid or base are added to the solution, buffers react with the added acid or base and resist large changes in pH.
Many biochemical reactions require specific buffers to work properly.
Therefore, making buffer solutions is an important laboratory skill in biochemistry.
For a brief review of how buffers work, let’s look at a generic buffer that consists of a weak acid, HA, and its conjugate base, A minus.
If an acid, H plus, is added to the buffer solution, the conjugate base A minus will react with it.
We can write the balanced equation for this acid-base neutralization reaction as A minus plus H plus gives HA.
If a base, OH minus, is added to the buffer solution, the weak acid HA will react with it.
We can write the balanced equation for this acid-base neutralization reaction as OH minus plus HA gives H2O plus A minus.
Therefore, a buffer resists changes in pH because it contains both an acid, HA, to neutralize added OH minus ions and a base, A minus, to neutralize added H plus ions.
There are two ways to make a buffer solution.
The first way to make a buffer solution starts with an aqueous solution of a weak acid, HA.
Since weak acids have a low percent ionization, in aqueous solution, most of the weak acid will stay as HA, and only a small percentage of HA will turn into A minus.
So, to represent an aqueous solution of a weak acid HA, in our particulate diagram, we start out with four HA molecules, and zero A minus anions.
Water molecules are left out of the particulate diagram for clarity.
Next, we add a solid salt, NaA.
When NaA dissolves in water, NaA turns into the Na plus ion and the A minus anion.
So if we add four particles of NaA to the aqueous solution of HA, they will dissociate and form four Na plus ions and four A minus anions.
Buffers work the best when the concentrations of the weak acid, HA, and its conjugate base, A minus are equal.
Since we have four HA molecules and four A minus anions in solution, we have a good buffer solution.
There are also four sodium ions present in the solution; however, sodium ions don’t react with water.
Since the sodium ions don’t react with water, they don’t affect the pH.
The second way to make a buffer solution also starts with an aqueous solution of a weak acid, HA.
However, this time, instead of adding a salt, we will add a strong base.
This time let’s start with 8 HA molecules in the particulate diagram.
Our next step is to add an aqueous solution of sodium hydroxide.
Sodium hydroxide is a strong base that dissociates into sodium ions, Na plus, and hydroxide ions, OH minus, when dissolved in water.
OH minus reacts with HA in an acid-base neutralization reaction to form H2O and A minus.
So, in the particulate diagram, if we pour in a solution that contains four particles of OH minus, the four OH minus ions will react with four HA molecules to form four A minus ions.
So, after the acid-base neutralization reaction is complete, we are left with four HA molecules and four A minus anions.
Since we have equal amounts of HA and A minus, we have a good buffer solution.
Since we added four particles of OH minus, we would also have four sodium ions in solution.
Notice how the buffer solution that we made the second way is identical to the buffer solution that we made the first way.
They both have four HA molecules, four A minus anions, and four Na plus ions.
Now that we know how to make buffers, let’s look at why buffers are most effective over a certain pH range.
Since buffers work the best when the concentration of HA and A minus are equal, let’s find the optimal pH for a buffer.
To find the pH of a buffer solution, we use the Henderson-Hasselbalch equation which says that the pH is equal to the pKa of the weak acid plus the log of the concentration of A minus divided by the concentration of HA.
Since the concentrations of A minus and HA are equal, the ratio of A minus to HA is one to one.
Since the log of one is equal to zero, the pH of the buffer solution is equal to the pKa of the weak acid.
If we add enough acid or base, eventually a buffer stops being effective.
This is often called “breaking a buffer.”
Let’s start with a buffer solution that has equal concentrations of HA and A minus.
If we add H plus to the buffer solution, the A minus reacts with the H plus to form HA.
Therefore, adding acid causes the concentration of A minus to decrease and the concentration of HA to increase.
When the ratio of A minus to HA becomes less than 1:10, buffer solutions “break” and lose their buffering action.
If we plug in a ratio of 1:10 in the Henderson-Hasselbalch equation, the log of 1 divided by 10 is negative one.
So the pH of the solution is equal to the pKa minus one.