I was told to design an experiment and run it during my chemistry lab period, and I came up with the following idea: Soil usually is limed with calcium carbonate containing products (agricultural lime) in order to raise the pH of the solution. Thus, I first isolated the amount of calcium from a soil sample made of known agricultural lime and pure soil (meaning no other special ingredients) in order to determine the amount of calcium carbonate given by the agricultural lime (room for error). Then I titrated a dilute soil sample with an automatic industrial titration apparatus with 0.1 M HCl and another solution of simply water and calcium carbonate of the same amount found in the same amount of soil in order to compare the pH curves (pH vs. titrant 0.1 M HCl). I found the pH curves for soil samples to reach the pH of pure HCl as the amount of titrant increased into excess (as I had expected) but the pH curves for the calcium carbonate and water solution indefinitely flat-lined at ~5 pH. Can anyone explain why? I know that calcium carbonate is highly insoluble in water, but when acid molecules "collide" with the calcium carbonate, it should react to form calcium chloride, water, and carbon dioxide. The purpose of my experiment was to see if the agricultural lime's manufacturer indeed gave the proposed amount of calcium carbonate, and to see if the the pH curves for pure calcium carbonate and soil were different in order to infer if the only buffering mechanism in soil is indeed calcium carbonate (highly unlikely, but it is a conclusion). I know that it sounds simplistic, but this is for a course before organic chemistry.