Does bicarbonation require water?

In an aqueous solution of $\ce{NaOH}$ exposed to the atmosphere one can probably expect some degree of $$\ce{NaOH(aq) + CO2(g) -> NaHCO3 (aq)}$$


But, what would happen if just a solid block of $\ce{NaOH(s)}$ was placed in some completely anhydrous environment, populated only by excess $\ce{CO2(g)}$ at STP? After all,


$$\ce{NaOH(s) + CO2(g) -> NaHCO3 (s)}$$


is technically a spontaneous process with $\Delta G^\circ = -76.91\ \mathrm{kJ/mol}$, but I can only think of three possibilities:


1. Nothing happens—no water (i.e., as a catalyst), no reaction!
2. A passivation layer of $\ce{NaHCO3(s)}$ forms on the surface, preventing the rest of the $\ce{NaOH(s)}$ inside from reacting^
3. No passivation layer—the $\ce{NaOH(s)}$ just keeps reacting away until it's all $\ce{NaHCO3(s)}$!




Rather 


$$\ce{2 NaOH(aq) + CO2(g) -> Na2CO3 (aq) + H2O}$$


and analogically on the solid $\ce{NaOH}$.


Bicarbonate in aqueous solutions cannot survive the excess of hydroxide, forming carbonate.


$$\ce{HCO3^{-}\ (aq) + OH- (aq)<=>> CO3^{2-}(aq) + H2O}$$


The same is further enforced by dehydration effect of solid hydroxide.


$$\ce{NaHCO3(s) + NaOH(s) -> Na2CO3(s) + H2O}$$


Only in excess of $\ce{CO2}$ Is formed bicarbonate:


$$\ce{CO3^{2-}(aq) + CO2(aq) + H2O <=>> 2 HCO3^- (aq) }$$


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