A casual look through the CRC Handbook of Chemistry and Physics will turn up dozens of substances that share with cerium(III) sulfate a preference for dissolving in cooler water. A few examples are:
sodium dihydrogen pyrophosphate hexahydrate
. . . and virtually any substance that's a gas at ordinary temperatures, including nitrogen, oxygen, hydrogen, helium, carbon dioxide, and ammonia.
This solubility behavior accounts for most of the noise your teapot makes before the water reaches its boiling point. Air that was dissolved in the cold water can't stay there as the temperature rises, and it bubbles loudly out of solution. By the time the water itself is hot enough to vaporize rapidly, most of the dissolved gases have left.
Table salt (sodium chloride) is almost indifferent to temperature in this regard: It's only 10 percent more soluble in boiling water than in ice water.
The thermodynamic "rule of thumb" is that if heat is given off when a substance dissolves, then that stuff will dissolve better in cold water than in warm. (This is a consequence of Le Chatelier's principle, discovered by the inventor of oxyacetylene welding.)
This solubility rule applies only to solids that are stable in the presence of water. Consider, for example, those instant "hot packs" that athletic trainers use, which produce heat by combining calcium chloride with water. Calcium chloride is unstable in the presence of water. The crystalline contents of the hot packs work because they are anhydrous (water free), having been roasted in a furnace. When mixed with water, most of the crystals do dissolve, releasing heat. Those that don't dissolve incorporate water molecules into their solid lattice, four to each calcium atom. The resulting crystals of calcium chloride tetrahydrate are stable in contact with water, but it turns out that they absorb heat upon dissolving, so the rule of thumb applies in reverse: solid CaCl2o¿4H2O dissolves better in hot water than in cold. And, of course, it would be useless in a hot pack.