Buffering agent

For an aqueous solution composed of both a weak acid or base and its conjugate, see Buffer solution. For uses not related to acid-base chemistry, see Buffer.

A buffering agent is a weak acid or base used to maintain the acidity (pH) of a solution near a chosen value after the addition of another acid or base. That is, the function of a buffering agent is to prevent a rapid change in pH when acids or bases are added to the solution. Buffering agents have variable properties—some are more soluble than others; some are acidic while others are basic. As pH managers, they are important in many chemical applications, including agriculture, food processing, biochemistry, medicine and photography.

Overview

A buffering agent can be either a weak acid or weak base. Buffering agents are usually added to water to form a buffer solution, which only slightly changes its pH in response to other acids and bases being combined with it, particularly a strong acid or a strong base. Another example is buffered aspirin which has a buffering agent, such as MgO or CaCO₃, that will help maintain the balance between the acid H-A (protonated) and the salt A⁻ (deprotonated) forms of aspirin as it passes through the acidic stomach of the patient. The H-A form of aspirin is a covalent molecule and is more absorbed by the stomach lining, leading to irritation. The antacid properties of these buffering agents help to maintain the balance toward the salt form by reducing the amount of stomach acid (HCl) which protonates the salt form, A⁻.

Mechanism of action

The way buffering agents work can be seen by calculating how little the pH of buffer solutions will change after addition of a strong acid or a strong base, whereas the same addition would significantly change the pH of a non-buffered solution. Using the Henderson-Hasselbalch equation we get an equilibrium expression between the acid and conjugate base in terms of the log of the ratio of the acid to conjugate base (the salt of the acid). The concentrations of the weak acid and its salt can change significantly, but the log of their ratio will not. The resulting pH of this combination can be found by using Le Chatelier's principle. For a simple numerical example, take the case where the concentrations of the weak acid and its salt are equal. If an added strong base halves the [HA] ([HA] decreases to [HA]-0.5[HA]), then [A⁻] will increase by [A⁻]+0.5[A⁻] since every molecule of HA that dissociates forms one molecule of [A⁻]. Thus the pH will be raised by a factor of log3 or 0.5 pH units (when the original [HA] and [A⁻] are equal):

$\textrm{pH} = \textrm{pK}_{a}+ \log \left ( \frac{[\textrm{A}^-]}{[\textrm{HA}]} \right )$

$\textrm{pH} = \textrm{pK}_{a}+ \log \left ( \frac{mol\textrm{A}^-}{mol\textrm{HA}} \right )$

$\textrm{pH}_{new} = \textrm{pK}_{a}+ \log \left ( \frac{mol\textrm{A}^-+0.5mol\textrm{A}^-}{mol\textrm{HA}-0.5mol\textrm{HA}} \right )$

$\textrm{pH}_{new} = \textrm{pK}_{a}+ \log \left ( \frac{\textrm{n}+0.5\textrm{n}}{\textrm{n}-0.5\textrm{n}} \right )$

$\textrm{pH}_{new} = \textrm{pK}_{a}+ \log \left ( \frac{1.5\textrm{n}}{0.5\textrm{n}} \right )$

$\textrm{pH}_{new} = \textrm{pK}_{a}+ \log \left ( \frac{[\textrm{A}^-]_{original}}{[\textrm{HA}]_{original}} \right )+ \log \left ( 3 \right )$

Thus, if originally the pK_a = 7 and [HA] = [A⁻], then the pH will be changed from 7.0 to 7.5 after addition of a strong base that halves [HA] and adds a half to [A⁻]. If a similar addition of 0.001 mol NaOH is made to 1L of non-buffered water, the pH would change from 7 to 11.

Examples

Agriculture

Monopotassium phosphate (MKP) is an example of a buffering agent. It has a mildly acidic reaction; when applied as a fertilizer with urea or diammonium phosphate, it minimizes pH fluctuations which can cause nitrogen loss and simultaneously provides potassium nutrient.

Humans

Buffering agents in humans, functioning in acid base homeostasis, are extracellular agents (e.g., bicarbonate, ammonia) as well as intracellular agents (including proteins and phosphate). A dilute mixture of monosodium phosphate and disodium phosphate can be made to approach physiological pH and can serve as an agent for phosphate loading or colonic lavage.

Paper

Calcium carbonate is frequently added to paper to ensure that it resists the effect of acids in the environment and does not prematurely age.

Shampoos and detergents

Most shampoos are buffered so that they are slightly acidic. The buffer counteracts the alkalinity of soap. If either of these substances become too acidic or too alkaline, they could damage fabric, and/or skin and hair.

See also

• Buffer solution
• Good's buffers

References

• nl:Buffer (scheikunde)