My prof told me that we can compare Lewis base strength based on electron donation to the $\ce{H+}$ proton - i.e sigma bond formation with the $\ce{H+}$ proton - and proceeded to conclude that $\ce{S^{2-}}$ is a weaker Lewis base than $\ce{O^{2-}}$, since the former is a weaker base than the latter. He then proceeded to generalize that the oxide ion is always a stronger Lewis base than the sulfide ion.

Is this a valid method of comparing the strength of Lewis bases? I have two issues:

  1. Isn't Lewis acid/base strength quantified using kinetics, not thermodynamics? Don't stronger LAs react faster, and weaker LAs react slower?

  2. Even if this comparison were valid, would the comparison still be valid outside of water solution?

Does this mean that this question is invalid? No medium is specified.

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I think he knows what he's talking about. In explaining the much greater stability of silver sulfide ($K_{\mathrm{sp}} \approx 6 \cdot 10^{-51}$) vs. $\ce{Ag2O}$ ($K_{\mathrm{sp}} = 3.5 \cdot 10^{-16}$), he notes that charge density does not explain this disparity. In fact, charge density would imply that the oxide salt should be more stable.

For instance, although $\ce{O^2-}$ is a stronger B/L base than $\ce{S^2-}$, $\ce{S^2-}$ can be (and usually is) a stronger Lewis base than $\ce{O^2-}$ since S has a lower EN than O. Furthermore, recall that Lewis base strength depends on the Lewis acid with with the Lewis base interacts.

The text goes on to mention other factors such as having a metal cation with a surfeit of valence d electrons, and how that can populate the empty valence 3d orbitals of $\ce{S^2-}$. (S has 3d orbitals?)

This makes sense; a good nucleophile has less stabilized electrons - i.e. electrons that are further from the nucleus. It still seems that the questions posted above put B/L acidity ahead as the main deciding factor for sigma-donor strength.


1) There is no, actually, a definitive, wide acknowledged measure of Lewis base/acid strength. For Brønsted acids/bases there is $\mathrm pK_\mathrm a$/$\mathrm pK_\mathrm b$ , which is thermodynamical (and actually depends on solvent, so there is no absolute scale of acid strength). [1]

Keeping what said above in mind, a qualitative measure of Lewis base strength would be its ability to donate its lone pair. Obviously, it grows with negative charge of the particle and lowering its electronegativity, there is nothing new with it. It should grow with shrinking of negative ion, also understandable. The most common Lewis acid is proton, so comparing strength of conjugated acid we can roughly estimate order of strength of Lewis base. [2]

2) Quantitatively - no. Qualitatively - yes, it would in most cases

[1] On a side not, in some cases there IS a detectable difference between thermodynamical and kinetic Brønsted acidity. For details see 'sterically hindered bases' in context of organic synthesis.

[2] This, however, comes with a little twist: a Lewis base working as a strong base for one element may work as a weaker base for another element. This depends on relative size of donating and receiving orbitals size/energy. Some food for thoughts may be found in solubility table, comparing solubility of, for example, Be, Ba, Hg and Pb salts.

  • $\begingroup$ Speaking of sterically hindered bases one might also be interested in "frustrated" Lewis acid/bases. $\endgroup$ – Dissenter Jun 22 '14 at 17:49
  • $\begingroup$ Speaking also of Lewis acid/base strength, can we kinetically quantify LA/LB strength? $\endgroup$ – Dissenter Jun 22 '14 at 17:50
  • $\begingroup$ @Dissenter highly doubt it. $\endgroup$ – permeakra Jun 24 '14 at 17:06

One of the keys in the definition of Lewis acids and Lewis bases is that they react to form a Lewis adduct. To compare the relative strengths of different Lewis bases, you would allow them to react with the same Lewis acid, form the required adduct and measure the enthalpy of formation of the adduct. The adduct with the largest enthalpy of formation would contain the strongest Lewis base.

1) So Lewis acid and base strength is about thermodynamics.

2) You could run the reaction in any solvent (or gas phase) as long as all of the Lewis acids or bases you want to compare are able to form an adduct in that solvent. There is no reason why the order of a series of enthalpies of bond formation has to remain constant when you move from solvent to solvent, so your relative strengths could well change from one solvent to another. The same applies to oxide \ sulphide relative Lewis base strengths if you substitute some other Lewis acid in place of the proton. Again the relative ordering might well change.

  • $\begingroup$ So the comparison would not be necessarily valid outside water solution? $\endgroup$ – Dissenter Jun 19 '14 at 18:31
  • $\begingroup$ The relative ordering of the comparison may turn out different when water is replaced with another solvent. $\endgroup$ – ron Jun 19 '14 at 18:33
  • $\begingroup$ That's odd because I've seen test questions regarding "strongest" electron donator and I'm fairly sure he expects us to rationalize electron donation strength through acid/base strengths. I've put an example question above in the OP. $\endgroup$ – Dissenter Jun 19 '14 at 19:01
  • $\begingroup$ What Lewis acid are they reacting with $\ce{BF_3 ,~ H^{+}}$, what solvent is this taking place in? I'm sure the standard assumption is proton, water, but it is poor practice not to specify them. Test questions are just another opportunity to teach. Calling out the solvent and the Lewis acid might make a student think - "ah, so there could be other choices". $\endgroup$ – ron Jun 19 '14 at 19:25
  • $\begingroup$ Wait, where do you see boron trifluoride? Also, it would be interesting if you had an example in which the strength of two Lewis acids/bases flipped depending on the solvent. I'll try to think of one myself too. I'd imagine that BF3 would be a good electrophile in water solution but a poorer electrophile in say a hydrocarbon solvent by nature of the solvent being less polar. $\endgroup$ – Dissenter Jun 19 '14 at 19:30

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