Buffer solution

Acids and Bases:
Acid-base reaction theories
pH
Self-ionization of water
Buffer solutions
Systematic naming
Redox reactions
Electrochemistry
Strong acids
Weak acids
Strong bases
Weak bases

Buffer solutions are solutions which resist change in pH upon addition of small amounts of acid or base. As a consequence of this definition, the theory of buffer solutions presented in this article uses the Arrhenius and Bronsted-Lowry notion of acids and bases, as opposed to the Lewis acid-base theory. See Acid-base reaction theories for further details on this subtle point. The scope of this article also does not consider buffer solutions prepared with solvents other than water.

Buffer solutions usually consist of either a weak acid and its salt or a weak base. The resistive action is the result of the equilibrium which is set up between the weak acid and the salt:

HA(aq) ↔ H⁺(aq) + A^-(aq)

(Where HA represents the weak acid, H⁺ the hydrogen ion component of the dissolved salt and A^- the anion component of the salt.) Two assumptions are made about the composition of this equilibrium:

1. All the A^- ions result from the salt. This is valid due to the acid's weakness - it supplies very little anions compared to the salt.
2. The HA acid remains unchanged. The high concentration of A^- and H⁺ ions means the equilibrium lies very much to the left.

If an alkali is added to the solution, hydronium ions mop it up. These ions are regenerated as the equilibrium moves to the right and some of the acid is broken down in to hydronium ions and anions. If an acid is added, the anions simply combine with the substance and once again pH is restored.

When writing about buffer systems they can be represented as salt/acid, or conjugate base/acid.

Contents

1 Applications 2 Illustration of a buffer solution in action 3 Illustration of the effect of buffer solutions on pH 4 Making buffer solutions 4.1 Citric acid-phosphate buffer

Applications

Their resistance to changes in pH makes buffer solutions very useful for chemical manufacturing and essential for many biochemical processes.

Buffer solutions are necessary to keep the right pH for enzymes in many organisms to work. A lot of enzymes work only under very precise conditions, if the pH strays too far out of the margin the enzymes slow or stop working and the organism dies. A buffer of carbonic acid (H₂CO₃) and bicarbonate (HCO₃^-) is present in blood plasma, to maintain a pH between 7.35 and 7.45.

Industrially buffer solutions are useful in fermentation processes, and for setting the correct conditions for the dyes used in colouring fabrics.

Illustration of a buffer solution in action

• <math>\mathrm{CH_3COONa/CH_3COOH:\,}<math>

Sodium acetate dissociates completely in water:

<math>\mathrm{CH_3COONa_{(s)} \to CH_3COO^-_{(aq)} + Na^+_{(aq)}}<math>

If an acid is added, the H⁺ ions will be consumed by the conjugate base in the buffer:

<math>\mathrm{CH_3COO^-_{(aq)} + H^+ \to CH_3COOH_{(aq)}}<math>

If a base is added, the OH^- ions will be neutralized by the acid in the buffer:

<math>\mathrm{CH_3COOH_{(aq)} + OH^- \to CH_3COO^-_{(aq)} + water}<math>

• <math>\mathrm{H_2PO_4^-/H_3PO_4}<math>
• <math>\mathrm{HPO4^{2-}/H_2PO_4^-}<math>
• <math>\mathrm{PO_4^{3-}/HPO_4^{2-}}<math>

Illustration of the effect of buffer solutions on pH

• <math>\mathrm{CH_3COONa/CH_3COOH:\,}<math>

The ionization constant is:

<math> \mathrm{K_a = \frac{[H^+][CH_3COO^-]}{[CH_3COOH]}}<math>

Since buffer solutions only involve weak acids and bases, it can be assumed that ionization of the acetic acid and hydrolysis of the acetate ions are negligible. Therefore, when the initial concentrations of the acid and conjugate base are the same, the pH of the buffer is equal to the pKa of the acid.

After the addition of HCl (a strong acid), complete ionization of HCl occurs:

<math> \mathrm{HCl_{(aq)} \to H^+_{(aq)}+Cl^-_{(aq)}}<math>

Neutralization of the HCl by acetate ions occurs:

<math> \mathrm{CH_3COO^-_{(aq)}+H^+_{(aq)} \to CH_3COOH_{(aq)}}<math>

The used up hydrogen ions change the number of moles of acetic acid and acetate ions, and hydrogen ions:

<math> \mathrm{moles\ of\ CH_3COO^- = initial\ moles\ of\ CH_3COO^- - initial\ moles\ of\ HCl}<math>

<math>\mathrm{moles\ of\ CH_3COOH = initial\ moles\ of\ CH_3COOH + initial\ moles\ of\ HCl}<math>

Now the formula for ionization constants can be used to determine the new [H⁺] and the new pH.

<math> \mathrm{K_a = \frac{[H^+][CH_3COO^-]}{[CH_3COOH]}}<math>

Making buffer solutions

Citric acid-phosphate buffer

Make up 0.1M citric acid and 0.2M phosphate solutions then mix as follows,

Citric acid-phosphate buffers

pH	0.2M Na2HPO4 /ml	0.1M Citric Acid /ml
3.0	20.55	79.45
4.0	38.55	61.45
5.0	51.50	48.50
6.0	63.15	36.85
7.0	82.35	17.65
8.0	97.25	2.75

de:Pufferlösung

nl:Buffer (scheikunde)