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Why is Ammonia stronger than Phosphine in terms of basic strength? I am confused as Nitrogen is more electronegative than Phosphorous.

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The strength of a Lewis base corresponds to the strength of its conjugate acid. As usual, a weak acid will give a stronger base, therefore parent acids with higher pKa's deprotonate to stronger Lewis bases. ammonium, $\ce{NH4+}$, has pKa=9.25, while phosphonium, $\ce{PH4+}$, is a much stronger acid with pKa=~-12. Basicity decreases down the group because proton affinity decreases as electronegativity decreases. The N-H bond is stronger than the P-H bond (hence $\ce{NH4+}$ is a weaker acid than $\ce{PH4+}$).

Lewis bases donate electrons from their HOMO, and for ammonia's MO, the HOMO, which corresponds to the lone pair (LP), is very similar in energy to the nitrogen AO, therefore the LP electron density can be considered to be more localized to the nitrogen. Presumably in a MO for an analogous species with a group 15 central atom lower on the periodic table (with lower electronegativity), the AO of the central atom would be higher in energy, so the HOMO electron density would be less concentrated on that atom.

Also, ammonia is more hybridized than phosphine. The s-p energy gap is greater for phosphorous, so the energy penalty of hybridization is higher, therefore the LP will have more pure s character than sp3 character, making it less directed.

Since the ammonia LP electron density is more concentrated on the central atom, and more directed, it makes it more nucleophilic.

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