At room temperature, about 10 g of NaHCO3 will dissolve in 100 mL of water. (Cleary, this solubility will alter should the temperature change, which is why I must specify a temperature.)
Now saturation is an equilibrium property, which we may represent by the following equation:
NaHCO3(s)⇌NaHCO3(aq). As with any equilibrium we write the the equilibrium equation:
Ksp=[NaHCO3(aq)][NaHCO3(s)].
As you know, the term [NaHCO3(s)] is meaningless; you cannot speak of the concentration of a solid; so it is removed from the equation:
Ksp=[NaHCO3(aq)]=[Na+(aq)][HCO3(aq)]. So if solid sodium carbonate is present, this equilibrium operates, sodium carbonate is saturated. If we added extra sodium salt (from whatever source), the ion product [Na+(aq)][HCO3(aq)] would be greater than Ksp and sodium bicarbonate would precipitate from solution.
Ksp, the solubility product, is tabulated for a host of sparingly soluble and insoluble salts. They must be measured.