Tricky exam question
The obvious choice for the source of carbonate is sodium bicarbonate (UPDATE: sodium bicarbonate is not soluble enough per other answer, but sodium carbonate is a possibility), which has a positive enthalpy of solution even at high concentrations, so let's call that B.

The acid is a little trickier. If you look up the standard enthalpy of solution of most acids, you'll find negative values (meaning the temp would go up when they are added to water). However, these values refer to infinite dilution in neutral water. Once you've added a little bit and the pH has dropped, the reaction $\ce{H+ + HO- -&gt; H2O}$, which has a large negative enthalpy, essentially stops happening and the enthalpy of solution becomes positive. This is easy to tell by looking at the maximum solubility as a function of temperature. If it increases, then the enthalpy change of solution at saturation is positive. So at the high concentrations in the problem (4 g per 30 mL), many acids that are solid at room temp will cause a temperature drop when dissolved. Citric acid is one example.

A solution of citric acid is certainly acidic enough that it will cause rapidly production of carbon dioxide gas when mixed with a solution of sodium carbonate.