The order of increasing acidity for the given compounds is as follows: Acetic acid < Butanoic acid < Hexanoic acid < Benzoic acid.
The acidity of a compound primarily depends on the stability of its corresponding conjugate base. The more stable the conjugate base, the stronger the acid.
Since all of the compounds are carboxylic acids, they all have the same acidic functional group (-COOH). However, the size and structure of the R-group attached to the carboxyl group determines the acidity.
Acetic acid has the smallest R-group (a methyl group), which does not stabilize the negative charge of the conjugate base effectively. Therefore, it is the weakest acid among the given compounds.
Butanoic acid has a slightly larger R-group (three carbons), and this increased size provides better electron delocalization, making it slightly more acidic than acetic acid.
Hexanoic acid has an even larger R-group (five carbons), allowing for increased electron delocalization and stability of the conjugate base. Thus, it is more acidic than both acetic acid and butanoic acid.
Finally, benzoic acid has an aromatic ring (phenyl group) attached to the carboxyl group. The delocalized pi electrons in the aromatic ring provide significant stabilization to the conjugate base, making benzoic acid the most acidic compound among the given compounds.