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$\underline{\text{Consider the following question:}}$

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The first difference I noticed in these compounds was the large number of oxygen atoms in (IV). This definitely means that there is a larger number of lone pairs in this compound. So the obvious conclusion is that compound (IV) is the most basic or the least acidic one. However none of the options seem to have (IV) as the least acidic compound.

Where have I gone wrong?

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    $\begingroup$ You should better consider that (I) is a diketone that can enolize. (II) is a lactone. (III) is an ether + ketone. (IV) is an ester + diketone. $\endgroup$
    – Maurice
    Commented Dec 31, 2020 at 18:05
  • $\begingroup$ @Maurice I understand. However I'm interested in knowing why this reasoning fails here opposed to it being the sole property in comparing the basicity of species like diamines and triamines $\endgroup$
    – newbie105
    Commented Dec 31, 2020 at 18:10
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    $\begingroup$ You have it all backwards. Basicity of a compound is unrelated to its acidity, and both are unrelated to the number of electron pairs. Say, hydrazine N2H4 has two of them, while ammonia has one; which of the two is a stronger base? $\endgroup$
    – Ivan Neretin
    Commented Dec 31, 2020 at 18:28
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    $\begingroup$ Related: chemistry.stackexchange.com/questions/128120/… $\endgroup$
    – Ivan Neretin
    Commented Dec 31, 2020 at 18:33
  • $\begingroup$ @IvanNeretin I see. There seems to be some intermingling of different definitions of base- Lewis base, Bronsted base. $\endgroup$
    – newbie105
    Commented Dec 31, 2020 at 18:34

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