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Which of the following is the most basic compound?

  1. $\ce{HS-}$
  2. $\ce{H3Si-}$
  3. $\ce{H2P-}$
  4. $\ce{Br-}$

I can't figure out which of these is the strongest base. Would it be $\ce{H3Si-}$, since $\ce{Si}$ is the least electronegative of all these central elements?

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  • $\begingroup$ The size of the atom carrying the charge is generally more important than the electronegativity. $\endgroup$ – Michael Lautman May 22 '19 at 15:29
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I agree with Doctor Insult's argument that "the stronger the base, weaker is the conjugate acid." In that sense, higher the $\mathrm{p}K_\mathrm{a}$ of acid, weaker is the acid, and hence, stronger is the conjugate base. The $\mathrm{p}K_\mathrm{a1}$ values of $\ce{HBr}$, $\ce{H2S}$, $\ce{H3P}$, and $\ce{H4Si}$ are -8.7, 6.89, 27, and ~35, respectively (Ref.1). Therefore, the order of the strength of their conjugate bases should be: $$\ce{H3Si- \gt H2P- \gt HS- \gt Br-}$$ Hence, strongest base is $\ce{H3Si-}$.

Note: Ironically, this is the same decreasing order of electronegativity of center atoms $\ce{Si, P, S, Br}$: $1.90 \lt 2.19 \lt 2.58 \lt 2.96$, respectively.

References:

  1. William L. Jolly, In Modern Inorganic Chemistry; 1st Edn.; McGraw-Hill, Inc.: New York, NY, 1984, p 177.
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Basic strength of a base is inversely proportional to the electronic negative charge density on the charged atom, because it will more readily to try to stabilize it charge by abstracting a proton. Size varies as Br>Si>S>P, so $\ce{H2P-}$ will be the strongest base. Even from the ${K_a}$ data, I concluded that $\ce{H2P-}$ will be the strongest base. Look up the ionisation constants of their conjugate acids. The stronger the base, weaker is the conjugate acid. Hope that helps.

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  • $\begingroup$ In this case, the charged atoms are all similar in size, so the electronegativity is the dominant effect. $\endgroup$ – Michael Lautman May 22 '19 at 16:59

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