In this article we will discuss about some chemical compounds like acid, base and buffers.
Acid and Base:
An acid may be defined as a substance which, when dissolved in water, produces hydrogen ions [H+] or in other words, it is the compound of electronegative element with ionisable hydrogen [H+]. Similarly a base or alkali is a substance which produces hydroxyl ions [OH–] when dissolved in water.
That is electropositive elements with ionisable hydroxyl groups [OH–] are the bases. A normal solution of any acid will exactly neutralise an equal volume of a normal solution of any base, no matter, how strong or weak, may be characteristics of two reactants. During neutralisation equal numbers of hydrogen ions of acid and hydroxyl ions of base unite to form water.
(i) Free Acidity:
Amount of acidity present in a solution in Free State and not combined with any other substance.
(ii) Actual Acidity:
Amount of [H+] ion concentration in a solution.
(iii) Total Acidity or Titratable Acidity:
Amount of free acid plus that present in combination. This can be titrated against a base.
The equilibrium governing the ionisation of water is as follows:
At a particular temperature the product of a number of hydrogen ions multiplied by a number of hydroxyl ions, is a constant.
Thus at 23°C.,
Concentration of [H+] and [OH–] ions in aqueous solution is such that their product is always equal to 10-14 gram-molecular weight per litre. This concentration product or dissociation constant of water can be denoted as Kw.
So the equation will be:
[H+] × [OH–] = Kw = 0.00000000000001 = 1 × 10-14.
In pure water the concentration of [H+] ions is equal to the concentration of [OH–].
So [H+] × [OH–] = 1 × 10-14 and [H+] = 1 × 10-7, [OH–] = 1 × 10-7.
Therefore pH of the pure water will be pH = log(1/[H+]) = log (1/[1 × 10-7]) = 7.0. So the pure water having equal number of [H+] and [OH–] ions is neutral at pH 7.0.
(cH) is a measure of extent of acidity or alkalinity of a solution. As the product of hydrogen- and hydroxyl- ion concentrations in a solution is constant, if the former is known, the latter can be calculated. Thus the hydrogen-ion concentration of neutral water is 10-7 and of normal sodium hydroxide is 10-14.
Buffers are usually a mixture of either weak acids with their salts of strong bases or strong acids with their salts of weak bases. The buffer system of the body has the capability of resisting the change in pH by accepting either the [H+] or [OH–] Ions. Buffering action is very important in the biological system in maintenance of normal acid- base balance of the body. It is most important in the action of enzymes, both in vivo and in vitro.
Henderson-Hasselbalch equation can be used for determination of pH of a buffer solution or for determination of relative concentration of the salt and acid required for achieving the normal pH.
The pH of the buffer solution prepared by mixing 35 ml of (N/10) acetic acid with 15 ml of (N/10) NaOH can be determined by the Henderson-Hasselbalch equation:
pH = pKa + log (salt/acid), where pKa is log (1/Ka). Ka is the dissociation constant of the acid.
When these two solutions are mixed up 15 ml of (N/10) NaOH will neutralise 15 ml of (N/10) acetic acid so as to form 15 ml of the salt, Na-acetate. So 20 ml of acetic acid will remain unneutralised.
So in the buffer solution, the salt and unneutralised acid ratio will be (15/20). The dissociation constant (Ka) of acetic acid is (1.86 × 10-5) and so the pKa will be log (1/Ka), that is 4.73.
So in Henderson-Hasselbalch equation;
pH = 4.73 + log (15/20)
= 4.73 + log 0.75
= 4.73 + 1.8751
If the ratio of salt and unneutralised acid ratio becomes 1:1, then the pH value of the buffer mixture will be equal to pKa of the acid.
As for example, if 30 ml of (N/10) acetic acid and 15 ml of (N/10) NaOH are mixed up then 15 ml of salt— Na-acetate will be formed and 15 ml of acetic acid will remain as unneutralised acid. So the salt and acid ratio will be (15/15)=1.
So interpolating the value in the Henderson-Hasselbalch equation the result will be:
pH = pKa + log (salt/acid)
= 4.73 + log (15/15)
= 4.73 + log 1
= 4.73 + 0
So the buffering action in this case is at its maximum and can react either as acid or as base.
Consider the equilibrium holding in a pure solution of a week acid is:
HA ⇋ [H+] + [A–]
By Law of Mass Action [H+] [A–]/[HA] = Ka or [H+] = Ka [HA]/[A–]
So log [H+] = log Ka + log [HA] – log [A–]
or, – log [H+] = – log Ka – log [HA] + log [A–]
According to Sorensen – log [H+] = pH, by similar reasoning – log Ka = pKa
Therefore, pH = pKa + log [A–]/[HA].