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Why is sodium carbonate less soluble in water than sodium bicarbonate? If you think about their structures, the only thing that is different is that sodium carbonate has two sodium atoms, while sodium bicarbonate has one sodium atom and one hydrogen atom instead of the second sodium atom. Sodium and hydrogen have different electronegativity values, giving sodium bicarbonate polar character.

Does the difference in solubility have to do with the dipole moment of the molecules? If so, how can it be explained in more detail?

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  • $\begingroup$ related chemistry.stackexchange.com/questions/5128/… $\endgroup$ – Mithoron May 8 '17 at 20:27
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    $\begingroup$ I suspect that it's a matter of the carbonate ion having a greater formal charge, causing it to be better solvated than the singly charged bicarbonate ion. Keep in mind that if I were certain of this, I would have written a formal answer ;) $\endgroup$ – airhuff May 8 '17 at 23:20
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Likely, this is due to hydrogen bonding. $\ce{NaHCO3}$ has one hydrogen bond donor site, but $\ce{Na2CO3}$ has none.

For an organic compound, one hydrogen bond for each four-five carbons may be enough to make it water soluble.

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According to the Born–Landé equation, the lattice energy is proportional to Z+Z-. Therefore in $\ce{Na2CO3}$ the Z+Z- value is 2 where as in $\ce{NaHCO3}$ the value of Z+Z- value is one. So the lattice energy is higher in $\ce{Na2CO3}$ than $\ce{NaHCO3}$. In addition to that the size of bicarbonate ion is higher than carbonate. Therefore packing efficiency will be higher in sodium carbonate than sodium bicarbonate. Thus more energy will require to disrupt $\ce{Na2CO3}$ than $\ce{NaHCO3}$. Therefore $\ce{Na2CO3}$ is less soluble in water than $\ce{NaHCO3}$.

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