Why are entries missing on a solubility data chart for ionic compounds?
In the case of magnesium chromate, good solubility in water has been known since 1931.  From Hill et al. [1], citing [2]:

The only figure on the solubility of the salt [as of 1940] appears to be that of Kohlraush[2], who reported that at 18°[C] 100 ml of saturated solution contained 60 g of magnesium chromate.

Most of the other compounds listed, as described in comments to the question, do not exist or are highly unstable because the ions are incompatible in the proposed solid. Iodide and sulfide ions are oxidized by iron(III), iodide is also oxidized by copper(II), chromate ions are reduced by iron(II) and tin(II), and tin(II) also reduces nitrate.  Carbonate ion is not redox-active, but fails to form a stable salt with some weakly-basic or amphotetic oxide forming metal ions like tin(II) or iron(III); a similar failure occurs with proposed silver hydroxide.
References
1.
Arthur E. Hill, Glenn C. Soth, and John E. Ricci, &quot;The Systems Magnesium Chromate—Water and Ammonium Chromate—Water from 0 to 75°&quot;, J. Am. Chem. Soc. 1940, 62, 8, 2131–2134. https://doi.org/10.1021/ja01865a059
2.
Joseph Mellor, Treatise on Inorganic and Theoretical Chemistry, Vol. II (London: Longmanns, 1931), p. 275.

Besides missing the reference on magnesium chromate, the chart has some errors associated with oxides and sulfides.
Magnesium oxide is listed as insoluble in water, which it is if it is calcined at a high temperature, but otherwise it becomes (kinetically) reactive forming the hydroxide; see this answer. Barium oxide is not &quot;soluble&quot;, but reacts in the same way as calcium, strontium, and low-temperature calcined magnesium oxides.
Among sulfides, the $\ce{S^{2-}}$ ion is a relatively strong base, so the alkali and alkaline earth metal sulfides that &quot;dissolve&quot; in water are in fact all prone to hydrolysis.  Calcium and magnesium sulfides do undergo more complete hydrolysis, as the corresponding hydroxides have limited solubility and their precipitation limits the basicity of the solution. Sulfide ion is also too strongly basic to coexist with ammonium ion, so ammonium sulfide as $\ce{(NH4)_2S}$ properly does not exist.