I know that acidity of an acid increases across a period (with electronegativity increase of the atom bonded to hydrogen) and hence $\ce{HF}$ is more acidic than $\ce{H2O}$ or $\ce{NH3}$. But the reason I learnt it is this way, is that difference in electronegativities of the atoms increases and hence bond becomes more polar and weaker (hence easier to donate a proton).

But shouldn't bond strength increase with difference in electronegativities? Also bond length of $\ce{HF}$ is shorter than that of $\ce{H2O}$ ($\pu{92 pm}$ vs $\pu{96 pm}$) and bond enthalpy of $\ce{HF}$ is $\pu{565kJ/mol}$ as opposed to $\pu{490kJ/mol}$ of water. Isn't this contradictory? Why is $\ce{HF}$ still more acidic than $\ce{H2O}$? Is it because the conjugate base of $\ce{HF}$ is highly solvated and more stable in an aqueous solution? Is there not a direct correlation between ease of bond dissociation and acidity?

P.S.: I understand the acidic trend down the group (stability of conjugate base increases and since the halogen atom becomes bigger, resulting in weaker bonds).

  • 1
    $\begingroup$ see this link . Did you get your answer? $\endgroup$ May 28 '20 at 12:31
  • $\begingroup$ Your welcome :-) $\endgroup$ May 30 '20 at 8:26

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