A Quick Guide to Buffers and pH

pH is the measure of acidity or basicity of a solution. Tightly maintaining the pH to the desired range is required in all experiments, and buffer solutions help with this. The scientific mechanism of how a buffer maintains pH is complex and will be the topic of this article.

A buffer, in ordinary terms, is something that protects or shields against an unwanted effect.

In bioscience, this unwanted effect is a drastic change in pH that might ruin your assay. And, one of the main functions of a buffer solution is to protect against pH fluctuations in the experimental medium.

The science behind how a buffer functions to maintain pH, however, is complex, and here we will learn about the theoretical concepts behind it, starting with a recap about our basic understanding of pH. And then we will discuss the mechanism buffers use to maintain the pH of a solution.

While understanding this concept, we will learn about an important term – pK, which as we will see is often used in deciding which buffer to use for your experiment.

Article table of contents:

The concept of pH

Significance of pK value of abuffer

References

The concept of pH

pH is a measure of the concentration of protons, or hydrogen ions (H⁺), in a solution. In a range of 0-14, the lower the pH, the more acidic a solution is, and the higher the pH, the more basic it is. A pH of 7 indicates neutrality.

Figure 1. Simple illustration of the pH scale where 7 is neutral, below 7 indicates an acidic pH, and above 7 indicates a basic pH.

Now let us get a little more technical. In chemistry, pH is mathematically defined as:

pH= - log₁₀[H⁺]

Here are a couple of things to note in general.

One, a square bracket ([ ]) indicates molar concentration. So, here [H⁺] means molar concentration of H⁺ ions in the solution.

Two, “p” (as in pH) of any value is the -log₁₀ of that value. Just as, the p in pH means -log₁₀[H].

The nice thing is, if you know the pH of a solution, you can derive the H⁺ ion concentration of that solution.

For example, let’s say the pH is 5. Using the equation pH= - log₁₀[H⁺], you’d have:

5 = -log₁₀[H⁺]

And you would then be able to get the value of [H⁺], or the H⁺ ion concentration of this solution from this equation.

With this in mind, let us see what would happen when an acid is in solution. For the sake of simplicity, since the structure of acetic acid (CH₃COOH) is less complicated, we will use acetic acid as our example:

Acetic acid dissociates by ionization in solution to hydrogen ions and acetate (CH₃COO^-). After ionization of an acid in solution, the chemical species formed is called the conjugate base.

So, in this case CH₃COO^- is the conjugate base for CH₃COOH. And acetic acid-acetate form a conjugate acid-base pair.

So, keeping this example of acetic acid in mind, this is what happens when an acid dissociates:

HA is the acid getting dissociated in solution. A^- is the conjugate base for this acid. The molar concentration of the conjugate base is [A^-] (recall brackets indicate concentration. And [HA] is the molar concentration of the acid.

Here are a couple of more examples, from inorganic acids very commonly used in labs.

Hydrochloric acid (HCl) and chloride ion (Cl^-) form the conjugate acid-base pair.

So now that we’ve covered what pH is, and how an acid dissociates in solution, here is another very important parameter, known as pK.

Every buffer has a characteristic pK, and whether a buffer is suitable for your experiments depends on its pK value. A buffer works best, when the pH it needs to maintain in the experiment, is very close to its pK value.

From one perspective, pK tells you how likely a buffer is to “give away” its hydrogen ion in a solution.

• If pK is less than 7, it means the substance is more likely to give away its H⁺ ions, so it's considered acidic. Strong acids have low pK values, while weak acids have pK values close to 7.
• If pK is greater than 7, it means the substance is less likely to give away its H⁺ ions, so it's considered basic. Strong bases have high pK values, and weak bases have pK values only slightly higher than 7.
• If pK is exactly 7, it's neutral, like water.

So, looking at the pK of a solution is a way for chemists to quantify and compare the strength of acids and bases.

The more important significance of pK is in the context of how a buffer works to maintain pH.

To understand this, let us take the example of acetic acid (CH₃COOH), and revisit how it dissociates in a solution of water:

For a reversible reaction like this, there is an equilibrium (or balance) between the two sides of the equation.

Depending on the conditions such as pH, the equilibrium may shift towards one of the two sides. For instance, in one case, we may have more CH₃COOH and less H⁺ + CH₃COO^-; that is, the equilibrium is shifted towards the left-hand side of the equation.

Conversely, there might be more H⁺ + CH₃COO^- and less CH₃COOH – with the equilibrium shifted more towards the right-hand side of the equation in this case.

As a third possibility, the two sides might be in a completely balanced state. Here the concentration of CH₃COOH and CH₃COO^- are equal. This is the most interesting scenario from the point of explaining what pK is.

pK of this acid is the pH at which it is half dissociated.

For CH₃COOH, the pH is 4.76. So, the pK of CH₃COOH is 4.76.

When acetic acid, at a pH of 4.76 is half-dissociated, the concentration of acetic acid (acidic form) and acetate ions (basic form) are roughly equal. This is the point where the acid is in a state of balance, or equal readiness for donating as well as for retaining its hydrogen ion.

And this equal readiness to donate or accept hydrogen ions is what allows the buffer to effectively resist changes in pH when small amounts of an acid or base are introduced, keeping the pH within a narrow range around the pK.

So, for an acetic acid/acetate buffer, it is most effective in maintaining the pH of the experimental environment at about 4.76.

Just in case you are interested, here are some mathematical details which might help in understanding what we discussed about pK and pH in a bit more depth and clarity.

This is what its mathematical definition is:

This equation, known as Henderson-Hasselbalch equation, is centrally important when we are trying to understand the concept of buffers. Here is why it is so important.

If the molar concentrations of the conjugated base and the acid are equal – that is, [ A^-] and [HA] are equal, we would have pH= pK + log 1, or pH=pK.

So, in other words, pK of an acid is the pH at which it is half dissociated, that is, [ A^-] = [HA].

With all of this in mind, let us now take a more detailed look into why pK value is so important in the context of a buffer.

Significance of pK value of abuffer

Each buffer has a characteristic pK. The tight range of pH that a buffer maintains is called the buffering capacity of that buffer, and its value is pK +/- 1.

The pK of the buffer you choose should be +/-1 of the pH that you need to maintain for your experiment.

For example, taking the instance of acetic acid that has a pK value of 4.76, this buffer will be effective in maintaining the pH in the range of 3.76-5.76 (that is, +/- 1 of 4.76).

Buffers help ensure that the pH of a solution is maintained within a narrow range even if other chemicals are added to the solution.

Because of this strict pH maintenance, experiments such as enzyme assays, protein purifications and cell cultures can be done properly in the lab. Now, let’s look closer at how a buffer maintains the pH of a solution and how the pK value is so relevant in that.

First, we need to understand what a weak acid is. A weak acid does not completely dissociate into ions in aqueous solution. And any weak acid, such as acetic acid, can act as a buffer.

To understand how, let’s take another look at the ionization equation of acetic acid, that we discussed earlier:

Imagine you start adding a base, that provides OH^- to this solution.

The OH^- ions of the base, react with the H⁺ ions of the dissociated acid, producing water (H₂O).

The graph below is known as titration curve. This curve depicts how the pH changes as more and more base (OH^- ions) is added to the acetic acid solution.

Figure 2. A titration curve demonstrating how a buffer composed of a solution of acetic acid maintains pH even when bases are added to it.

Please note that for a pH around 4.76, which is the pK of acetic acid, a large amount of OH^- ions are needed to raise the pH of the solution even modestly. On the other hand, a large amount of H⁺ ions (if an acid is added to this solution) are needed to even slightly lower the pH.

This range, which amounts to pK+/- 1, is the buffering capacity of this acetic acid buffer, and ensures that the pH of the solution remains within 4.76 +/- 1.

This is the chemical mechanism behind how a buffer maintains the pH of solution.

In light of this concept, we now come to the most important question: how do I choose the buffer needed for my experiment based on pK value?

You know the pH range needed for your experiment, and the pK of commercial buffers are provided by the manufacturer.

The pK of the buffer you choose should be +/- 1 of the pH that you need to maintain for your experiment.

As a validation of this concept, think of what goes on in biological fluids – most of which need a pH of around 7 to be maintained. The natural buffers in those fluids are ions such as bicarbonate (HCO₃^-) that has a pK of 6.35, and phosphate (H2PO₄^-) with a pK of 6.82. If you notice, the pK of these buffering chemicals are 7 +/- 1.

So, for your experiments, we have a really helpful chart that shows different GoldBio buffers along with their pH ranges and pK values that you might want to check out.

References

Nelson, David L. (David Lee), 1942-. (2005). Lehninger principles of biochemistry. New York :W.H. Freeman