O + in the solution comes mainly from the first dissociation step. . . Factors Affecting Acid Strength Having discussed quantitatively the strengths of acids and bases, we come to a stage where we can calculate the pH of a given acid solution.
But, the curiosity rises about why should some acids be stronger than others? What factors are responsible for making them stronger? The answer lies in its being a complex phenomenon. But, broadly speaking we can say that the extent of dissociation of an acid depends on the strength and polarity of the H-A bond.
of common ion effect . It can be defined as a shift in equilibrium on adding a substance that provides more of an ionic species already present in the dissociation equilibrium. Thus, we can say that common ion effect is a phenomenon based on the Le Chatelier’s principle discussed in section . .
In order to evaluate the pH of the solution resulting on addition of .05M acetate ion to .05M acetic acid solution, we shall consider the acetic acid dissociation equilibrium once again, HAc(aq) H + (aq) + Ac – (aq) Initial concentration (M) . . Let x be the extent of ionization of acetic acid. Change in concentration (M) –x +x +x Equilibrium concentration (M) .
-x x . +x Therefore, K a = [H + ][Ac – ]/[H Ac] = {( . +x)(x)}/( . -x) As K a is small for a very weak acid, x<< .
Problem . Calculate the pH of a .10M ammonia solution. Calculate the pH after . mL of this solution is treated with .
mL of .10M HCl. The dissociation constant of ammonia, K b = . × – NH + H O → NH + + OH – K b = [NH + ][OH