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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,
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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.
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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.
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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.
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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.
