Why is the solubility of sodium acetate trihydrate so much lower than anhydrous sodium acetate? According to Wikipedia:

Solubility of anhydrous sodium acetate in water at 0 $^{\circ}$C is 119 g/100 mL (AKA 1.495 mol/100mL)

Solubility of trihydrate sodium acetate in water at 0 $^{\circ}$C is 36.2 g/100 mL.

This makes no sense to me because it should take MORE of the trihydrate sodium acetate (about 1.495 * 136.08 g) to dissolve the same amount of sodium acetate.

So this begs two questions:

1.) Is my understanding of hydrates incorrect? Are their bonds somehow harder to solvate than the anhydrous version? I think this can't be correct because the melting point of sodium acetate trihydrate is 58 $^{\circ}$C compared to the anhydrous melting point of 324 $^{\circ}$C. 

2.) Is my understanding of solubility incorrect? I thought solubility is a measure of the quantity of solute that could be dissolved in a solvent. In other words, it seems that MORE sodium cations and acetate anions can be dissolved in 100 mL of water if they came from an anhydrous crystal rather than a trihydrate crystal...
How can that be true?
The values of solubility for these salts are $100\%$ correct. Your understanding of solubility is also correct. What you didn't know (probably) is that solutions can exist in a non-equilibrium state. In case of anhydrous sodium acetate, it dissolves in water forming so-called supersaturated solution, which is metastable. Eventually, crystals of sodium acetate trihydrate will precipitate out of the solution, forming a system which is in the thermodynamic equilibrium.