Of the minerals you cite, clays and micas will act as bases in reaction with acidity. Quartz will also, but at very slow rates. Titanite - probably not much.

A very good resource, if you can locate it, is Reviews in Mineralogy, Vol. 31 (Mineralogical Society of America), edited by A.F. White and S.L. Brantley, Chemical Weathering Rates of Silicate Minerals (1995).

If the neutralization is happening very quickly, one might wonder whether, amongst the clays, there may be interstitial carbonates. Or perhaps you are titrating in small fluxes of acidity to a system with very high available surface area for reaction (especially the clays).

One more thing according chemical test of rock, in my rock it has Fe2O3, Mn, Cu, Zn, Hg, and As, which one is it as acid, because at first my AMD is pretty acid, its pH around 4 now it is going to reach around 8 in just 2 months, do you have any idea what is really happening? According to my thin section test and my chemical test of rock, it is kind confusing to me.

The identification of the iron in your rock as "Fe2O3" is just a convention derived from igneous petrology. These are fictional components - you can compute the actual concentration of Fe (to compare with, for example Zn) by converting the percentage value of Fe2O3 to mole % Fe2O3, recognizing there are 2 moles of Fe for every mole of Fe2O3, and then calculating the concentration of Fe by applying the atomic weight of Fe (and adjusting the units to be in mg/kg.

Because you seem to have an acid-0generating material, we can probably assume that much of the Fe and most/all of the Cu, Zn, Hg and As are present in the rock initially as sulfide minerals, for example chalcopyrite (CuFeS2), sphalerite (ZnS). If you have thin sections, you can identify these specifically for your rock; if you need more help for small mineral concentrations, see if you can get access to an X-ray Diffraction apparatus. The metals in their original sulfide form are present neither as acid nor bases, in the sense that they are little or not at all affected by hydrolysis (reaction with H+). However, as you know, they will dissolve during oxidation and produce dissolved metal (and sulfate) ions in solution, and the concentration that can remain in solution will be a function of the pH of the bulk solution.

There is an important conundrum to face in dealing with such minerals: the pH of the system is a *response* variable to the acid-base balance. However, at each stage of the solution's evolution, pH also will be a *master* variable that controls the solubility of the transition metals in solution. So it is NOT a linear system; there are feedbacks that you need to consider, and the role of the non-sulfide gangue (clays, carbonates, micas) at neutralizing acidity will determine how the pH evolves, and th3erefore, at least providing enough time for reactions to proceed, what metals concentrations you can expect to see.

If you are working in an experimental system, it is very important for you to consider how your experiment relates to the environment you actually wish to assess in the natural world.

There are many very bright and experienced geochemists on this group: I am sure some of them can provide you great insights, too.

Apart from the oxidation of sulphide minerals and dissolution of acid consuming minerals, pH will also be controlled by the precipitation and dissolution of secondary minerals. In order to be sure, have a look at the secondary minerals that are likely to form in your system by carrying out saturation indices calculations based on the solution chemistry. You can use common tools such as PHREEQC or Visual Minteq. This will give you an idea of the minerals that are likely to form and influence your pH besides the primary minerals. It is also important to note that even slow dissolving minerals can dissolve at a significantly fast rate at low pH thereby becoming potential acid consuming minerals.