Buffers

In some situations, a change in pH can cause extreme reactions to occur. For example, the pH of blood should be from 7.38-7.42. If it changes to 7 or 8, we would die. We must therefore have a mechanism to protect against even a slight change in pH. The chemical that does this is called a buffer. Buffers maintain pH, they do not make solutions neutral.

Buffer Components:

Buffers are made of two solutes, a conjugate pair containing a weak acid and a weak base.

Here are some examples:

HC₂H₃O₂ and NaC₂H₃O₂

H₂CO₃ and HCO₃ ^-

NH₄ ⁺ and NH₃

pH of a Suffered Solution

The pH of a buffer depends on

k_a (or k_b for a base)
the ratio of [A^-]/[HA]

Diluting a buffer does not change its pH because it does not change the ratio or the k_a.

Follow along with this video to learn how to calculate the pH of a buffer.

Note that you have a choice of what equation to use to begin a problem like th one above. You can either write the k_a equation using the acid or the k_b equation using the base.

There is a shortcut to calculate the pH of a buffer solution. The basic k_a expression can be rearranged to another very useful equation known as Henderson-Hasselbalch equation.

Henderson- Hasselbalch

pH = pk +log [base]A/[acid] H

This equation can be used for buffers with the initial concentrations of the buffer components. This equation can be written for pOH as well, but it is not used as commonly.

Let's do the same problem done in the video above. But this time we will use the Henderson-Hasselbalch equation.

What is the pH of a buffer solution containing 0.90 M HC2H3O2 and 1.10 M NaC2H3O2?
pH = pk +log [A⁻]/[HA]
pH = -log(1.8 x 10⁻⁵) + log (1.10/0.90)
pH 4.74 + 0.08715
pH = 4.83

How Buffers Work

Now that you know how to calculate the pH of a prepared buffer, let's talk about how buffers actually work and how to prepare them. A buffer works by neutralizing any strong acid or base added. This allows the original pH to remain constant or to change only by a negligible amount. A buffer solution contains a large concentration of both members in a conjugate acid-base pair. The conjugate acid in the buffer will react to neutralize added base and the conjugate base in the buffer will react to neutralize added acid.

Preparing a Buffer of a Given pH

Buffers are usually made so that the ratio of [A^-]/[HA] is near 1. Look at the Henderson-Hasselbalch equation below. What do you notice when the ratio of [A^-]/[HA] is exactly 1?

$LaTeX: pH=pk_a+log \frac{[A^-]}{[HA]}$

When this ratio is exactly 1, pH = pk_a. This is true any time [acid] = [base].

Buffers are also made to work around a certain pH. Here is how a buffer of a specified pH is prepared:

Choose a weak acid with a pk_a near the desired pH. The general rule is that the desired pH = pk_a ± 1.
Adjust the ratio of weak base to weak acid ([A^-]/[HA]) to get the exact desired pH.

By comparing the pH of a solution to the pKa of the acid in the solution, the concentration ratio between the acid and base forms of that acid (sometimes called the protonation state) can be determined. For example, if pH < pKa, the acid form has a higher concentration than the base form. If pH > pKa, the base form has a higher concentration than the acid form. Applications of this relationship include the use of acid-base indicators, the protonation state of protein side chains (including acids or proteins with multiple labile protons), and the pH required for acid-catalyzed reactions in organic chemistry.

Buffer Capacity

There is a limit to the amount of strong base or strong acid that a buffer can neutralize. This limit is known as buffer capacity. Buffer capacity is the number of moles of base (or acid) that, when added to 1 liter of buffer solution, changes its pH by 1 unit. Buffer capacity is determined by the actual concentration (or number of moles) of each component present NOT by their ratio.

The more A^- and HA molecules available, the less of an effect the addition of a strong acid or strong base will have on the pH of a system. Consider the addition of a strong acid such as HCl. Initially, the HCl donates its proton (H⁺) the weak base (A^-) through the reaction

A^- + HCl → HA + Cl^-

This changes the pH by lowering the ratio of [A^-]/[HA], but as long as there is still a lot of A^- present, the change in pH will be small. But if we keep adding HCl, the weak base A^- will eventually run out. Once the A^- is gone, any additional HCl will donate its proton to water (HCl + H₂O ⇌ H₃O⁺ + Cl^-). This will dramatically increase the [H⁺] and so the pH drops.

Remember to work on the module practice problems as you complete each section of content.

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