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Why is $\ce{Ph-CH2-COOH}$ more acidic than $\ce{CH3COOH}$ although the equilibrium of ionization lies mainly backward for both of them?
Attempt:
In the ionized form of $\ce{Ph-CH2-COOH}$ we have a $\ce{Ph-CH2}$-group donating its electron to carbonyl carbon while in case of acetic acid we have a methyl group donating its electron to the carbonyl carbon.
But which one of them is stabilized more?
The $\mathrm pK_\mathrm a$ difference between phenylacetic acid and acetic acid is around 0.5 and the trends become more obvious when other arylalkanoic acids are included in the comparison. Numbers in red are the $\mathrm pK_\mathrm a$ values.
Apparently, phenylacetic acid is a stronger acid than acetic acid. Adding a further phenyl subsituent even increases the acid strength.
The effect levels off with the distance of the phenyl substituent and the carboxylate. Note that 4-phenylbutyric acid is just as strong as acetic acid. It's as if there isn't any phenyl group at all.
Would you think that there is a $-I$ effect of the phenyl substituent on the $\ce{-COOH}$ group?
