16. Ions and Equilibrium;
       Acids and Bases
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       The difference between strong and weak acids

Whether an acid will behave as a strong or weak acid thus depends both on the acid and the solvent. Because the fluoride ion, F-, is small, thus permitting the proton to come close to it and feel a strong electrostatic attraction, it is a moderately strong B-L base.

Even though water molecules are present in great excess, the attraction of F- for protons is great enough that HF is only partially dissociated in aqueous solution. In contrast, the Cl- ion is large, with a diffuse electron density, and does not permit the proton to approach as closely. It attracts protons more weakly, and is a weak B-L base.

So many molecules of H2O surround each Cl- ion that they overwhelm Cl- and compete successfully for the available H+ ions, thereby pushing the dissociation equilibrium effectively to completion. Hence HCl is classed as a strong acid in water.

In methanol, Cl- ions in similar concentration find no difficulty in competing successfully with an excess of CH3OH molecules for the available H+, because of the very small attraction of CH3OH molecules for H+. Even in methanol, however, perchloric acid is a strong acid because the perchlorate ion, ClO4-, has less attraction for protons than methanol molecules do. HClO4 is the strongest of all the common acids because ClO4- is the weakest of all the B-L bases.



Sulfuric acid can lose two protons. The first dissociation is that of a strong acid, and is complete in aqueous solution:

The bisulfate ion, HSO4-, is more reluctant to lose another positive ion, since it already has one negative charge. The sulfate ion is a strong B-L base, and competes successfully with water molecules for the proton. Therefore HSO4- is a weak acid with a measurable dissociation constant:

Phosphoric acid has three dissociating protons, of varying degrees of weakness:


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