Buffer solution

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Template:Acids and Bases Buffer solutions are solutions which resist change in hydronium ion concentration (and consequent pH) upon addition of small amounts of acid or base, or upon dilution. Buffer solutions consist of a weak acid and its conjugate base. The resistive action is the result of the equilibrium between the weak acid (HA) and its conjugate base (A^-):

HA(aq) + H_{2</sup>O(l) ⇌ H3O⁺(aq) + A^-(aq)}

Any alkali added to the solution is consumed by hydronium ions. These ions are mostly regenerated as the equilibrium moves to the right and some of the acid dissociates into hydronium ions and the conjugate base. If a strong acid is added, the conjugate base is protonated, and the pH is almost entirely restored. This is an example of Le Chatelier's principle and the common ion effect. This contrasts with solutions of strong acids or strong bases, where any additional strong acid or base can greatly change the pH.

When writing about buffer systems they can be represented as salt of conjugate base/acid, or base/salt of conjugate acid. It should be noted that here buffer solutions are presented in terms of the Brønsted-Lowry notion of acids and bases, as opposed to the Lewis acid-base theory (see acid-base reaction theories). Omitted here are buffer solutions prepared with solvents other than water.

Calculating pH of a buffer

The equilibrium above has the following acid dissociation constant:

<math> \mathrm{K_a = \frac{[H^+][A^-]}{[HA]}}</math>

Simple manipulation with logarithms gives the Henderson-Hasselbalch equation, which describe pH in terms of pK_a:

<math>pH=pK_a+log_{10}\frac{[A^-]}{[HA]}</math>

In this equation

1. [A-] is the concentration of the conjugate base. This may be considered as coming completely from the salt, since the acid supplies relatively few anions compared to the salt.
2. [HA] is the concentration of the acid. It again is minimally dissociated, and hence can be considered a constant.

Illustration of Buffering Effect: Sodium acetate/Acetic acid

The acid dissociation constant for acetic acid-sodium acetate is given by the equation:

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

Since this equilibrium only involves a weak acid and base, it can be assumed that ionization of the acetic acid and hydrolysis of the acetate ions are negligible. In a buffer consisting of equal amounts of acetic acid and sodium acetate, the equilibrium equation simplifies to

<math> \mathrm{K_a = [H^+]}</math>,

and the pH of the buffer as is equal to the pK_a.

To determine the effect of addition of a strong acid such as HCl, the following mathematics would provide the new pH. Since HCl is (a strong acid) it is completely ionized in solution, and increases the concentration of H⁺ in solution, which then neutralizes the acetate by the following equation.

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

The consumed hydrogen ions change the effective number of moles of acetic acid and acetate 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>

After accounting for volume change to give concentrations, the new pH could be calculated from the Henderson-Hasselbalch equation. Any neutralization will result in a small change in pH since it is on a logarithmic scale.

Applications

Their resistance to changes in pH makes buffer solutions very useful for chemical manufacturing and essential for many biochemical processes. The ideal buffer for a particular pH has a pK_a equal to the pH desired, since a solution of this buffer would contain equal amounts of acid and base, and be in the middle of the range of buffering capacity.

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 enzyme can denature, thus permanently disabling its catalytic activity. 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.

Also, buffer solution can be used in chemical analysis and calibration of pH meters.

Common Buffer Compounds used in Biology

** Values are approximate

More buffers at Sigma-Aldrich

Making buffer solutions

Citric acid-phosphate buffer

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

See also

• Buffering agent
• common ion effect

External links

• Example buffer recipes
• Derivation and discussion of Henderson-Hasselbalch equationbg:Буфер (разтвор)

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