In this article we will discuss about:- 1. Definition of pH 2. Variation in pH and Interpretation 3. Determination.
Definition of pH:
pH is defined as the negative of logarithm (base 10) of the hydrogen ion concentration.
Or,
It is defined as the logarithm of the reciprocal of the hydrogen ion concentration.
All vital activities are affected by H+ concentration. Hydrogen ion concentration must be ascertained before the pH is calculated. For strong electrolytes, [H+] may be substantially the same as the total concentration, if complete ionisation is assumed. But for weak electrolytes, [H+] must be obtained by the calculation from the ionisation constant.
Variation in pH and Interpretation:
A solution of pH 3 contains 10-3 gram H+ per litre.
A solution of pH 5 contains 10-5 gram H+ per litre.
A solution of pH 8 contains 10″8 gram H+ per litre.
A solution of pH 3 has 10 times the [H+] of one of pH 4 and 100 times that of a solution of pH 5.
As [H+] increases, pH decreases in such a way that for one unit increase in pH the [H+] increases 10 times. It means that the higher the pH the lower will be the acidity.
A neutral solution has pH 7. Pure water (the neutral solution) is ionised as follows:
[H+] x [OH−] = Kw
or, [H+] x [OH−] = 1 x 10−14
Putting logarithm on both sides,
log [H+] + log [OH−] = log1 + log10−14.
or, −log[H+] – log[OH−] = 0 +14 log 10. [Multiplying both sides by – sign]
or, pH + pOH= 14 [Since, pH = 7].
... pH = pOH i.e. [H+] = [OH−].
Therefore, a solution with pH less than 7 is acid and higher than 7 is alkaline. The pH range is 0 to 14 only. It should be noted that the pH scale is logarithmic, not numerical.
pH 6.5 does not represent a [H+] half-way between 6 and 7.
Actually, pH 6.5 = [H+] 3.2 x 10−7
pH 6.0 = [H+] 10 x 10−7 = [H+] 10−7
Determination of pH:
For weak electrolytes with which the physiologic chemistry is concerned, this may be calculated by the law of mass action:
HA = Un-dissociated weak acid, Ka = Dissociation constant for the acid.