The acid-base titrations involve the determination of the acidity or alkalinity of the solutions and the purity of carbonates and alkaline earth oxides.
For the acid-base titration, the equivalence point is characterized by a pH level that is a function of the acid-base strengths and the concentrations of the analyte and the titrant.
Written by: J.A.D.I.Thathsarani, BSc.(Hons) special in Chemistry, University of Kelaniya (UOK)
However, the pH of the endpoint may or may not correspond to the pH at the equivalence point.
If both the acid and base are strong electrolytes, the solution at the endpoint will be neutral and have a pH of 7, but if either the acid or base is a weak electrolyte, the salt will be hydrolyzed to a certain degree and the solution at the equivalence point will be slightly alkaline or acid.
The exact pH of the solution at the equivalence point can readily be calculated from the ionization constant of the weak acid or base and the concentration of the solution.
For any actual titration, the correct endpoint will be characterized by a definite value of the hydrogen ion concentration [H3O+] of the solution, the value depending upon the nature of the acid and the base.
The acid-base indicators change colour according to the hydrogen ion concentration of the solution. The main feature of these indicators is that the change from a predominantly 'acid' to a predominantly 'base' takes place within a small interval of pH (colour change interval).
The position of the colour change interval in the pH scale varies widely with different indicators.
It should be selected as an indicator that exhibits a distinct colour change at the pH close to the corresponding to the equivalence point.
The mechanism of the neutralization process can be understood by studying the changes in the hydrogen ion concentration during the appropriate titration. The change in pH in the neighbourhood of the equivalence point is of the greatest importance, as it enables an indicator to be selected which will give the smallest titration error.
The curve obtained by plotting pH as the ordinate against the percentage of the acid neutralized (or the volume of alkali added) as abscissa is known as the neutralization curve.
In the titration of any strong acid with a strong base, there are three regions of the titration curve that require different kinds of calculations: For example, let's consider the titration of 50.0 mL of 0.100 M HCl with 0.2 M NaOH.
Where the subscript 'a' indicates the acid (HCl) and the subscript 'b' indicates the base (NaOH).
For the reaction of a strong base with a strong acid, the only important equilibrium reaction is as below.
Before the equivalence point, HCl is present in excess and the pH is determined by the concentration of excess HCl. Initially, the solution is 0.100 M in HCl, which since HCl is a strong acid, means that the pH is calculated as below.
The equilibrium constant for the reaction 1.1 is Kw = 1.00 x 10-14 . Since this is such a large value we can treat reaction 1.1 as though it goes to completion. After adding 10.00 mL of NaOH, therefore, the concentration of excess HCl is calculated as below.
So, pH value of solution is 1.3
At the equivalence point,
The volume of NaOH needed to reach the equivalence point is calculated as below.
At the equivalence point, the moles of HCl and the moles of NaOH are equal. Since neither the acid nor the base is in excess, the pH is determined by the dissociation of water.
Thus, the pH at the equivalence point is 7.00
Finally, for volumes of NaOH greater than the equivalence point volume, the pH is determined by the concentration of excess OH-. For example, after adding 30.0 mL of titrant the concentration of OH- is calculated as below.
To find the concentration of H3O+, we use Kw expression.
According to the concentration of H3O+ ions, pH is 12.10
If the acid is titrant, the curve is a mirror image as above.
**The size of the vertical region will depend on the concentration of the acid and the base.
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