Why does silver sulfate not precipitate in this double displacement reaction? Aqueous silver nitrate is mixed with aqueous sodium sulfate.  We were asked to perform this reaction in lab in my chemistry class, and we did not observe a precipitant being formed. However, the equation indicates otherwise:

$$\ce{2AgNO3 (aq) + Na2SO4 (aq) -&gt; Ag2SO4 (s) + 2NaNO3 (aq)}$$

This shows the formation of $\ce{Ag2SO4}$, silver sulfate. I found in the Wikipedia article for silver sulfate a reference to "ruby red illumination" and thus am led to believe that, like some silver-based substances used in photographic film, the precipitant might have been immediately been affected by light in the room and thus decomposed into something soluble in water. I have not, however, been able to find references to this anywhere else; I have found here a reference to it being classified as "sparingly soluble," but I haven't the slightest idea what this means.
Since the lab is past due now, I'll give what I think is the answer.

$\ce{Ag2SO4}$ is somewhat soluble in water. It is most likely that you simply didn't get enough to cause precipitation. 

The other possibility, which I don't think applies here, is that you have a supersaturated solution. There are some precipitates for which crystals are just stubborn to nucleate.