Buffer definition and examples in chemistry
In chemistry buffer definition and examples.It is a solution containing either a weak acid and its salt or a weak base and its salt , which resists changes in pH . In other words, a buffer is an aqueous solution of a weak acid and its conjugate base or a weak base and its conjugate acid.
Buffers are used to maintain a stable pH in a solution because they can neutralize small amounts of additional base acid.
For a given buffer solution, there is a working pH range and a defined amount of acid or base that can be neutralized before the pH changes. The amount of acid or base that can be added to a buffer before changing its pH is called its buffer capacity.
The Henderson-Hasselbalch equation can be used to measure the approximate pH of a buffer. To use the equation, the initial concentration or stoichiometric concentration is entered in place of the equilibrium concentration.
The general form of a chemical buffer reaction is:
HA ⇌ H + + A –
Also known as: Buffers are also called hydrogen ion buffers or pH buffers.
Examples of Buffers
- blood – contains a bicarbonate buffer system
- TRIS buffer
- phosphate buffer
Buffers pH range :
As noted, buffers are useful over specific pH ranges. For example, here is the pH range of common buffering agents:
|citric acid||3.13., 4.76, 6.40||2.1 to 7.4|
|acetic acid||4.8||3.8 to 5.8|
|KH 2 PO 4||7.2||6.2 to 8.2|
|borate||9.24||8:25 a.m. to 10:25 a.m.|
|CHES||9.3||From 8.3 to 10.3|
When a buffer solution is prepared, the pH of the solution is adjusted to bring it into the correct effective range. Typically, a strong acid, such as hydrochloric acid (HCl) is added to lower the pH of the acid buffers. A strong base, such as a solution of sodium hydroxide (NaOH), is added to raise the pH of the alkaline buffers.
How buffers work?
In order to understand how a buffer works, let’s take the example of a buffer solution obtained by dissolving sodium acetate in acetic acid. Acetic acid is (as we can say from the name) an acid: CH 3 COOH, while sodium acetate dissociates in solution to give the conjugate base, the acetate ions of CH 3 COO – . The equation for the reaction is:
CH 3 COOH (aq) + OH – (aq) ⇆ CH 3 COO – (aq) + H 2 O (aq)
If a strong acid is added to this solution, the acetate ion neutralizes it:
CH 3 COO – (aq) + H + (aq) ⇆ CH 3 COOH (aq)
This shifts the balance of the initial buffer reaction, keeping the pH stable. On the other hand, a strong base would react with acetic acid.
Most buffers operate over a relatively narrow pH range. An exception is citric acid because it has three pKa values. When a compound has multiple pKa values, a larger pH range becomes available for a buffer. It is also possible to combine buffers, provided that their pKa values are close (differing by 2 or less), and to adjust the pH with a strong base or an acid to reach the required range. For example, McIvaine’s buffer is prepared by combining mixtures of Na2PO4 and citric acid. Depending on the ratio between the compounds, the buffer can be effective from pH 3.0 to 8.0.
A mixture of citric acid, boric acid, monopotassium phosphate and diethylbarbituic acid can cover the pH range from 2.6 to 12!
you may also look at
You may also look at