[Loler]

Hello all,

I am working on my senior year project, and I am running into some problems trying to interpret some of my data.

I collected some groundwater samples and analyzed them for major ions and cations. The groundwater samples are from a wetland, that has two different layers. The top layer is organic soil, and the bottom layer is mineral. The water is moving from the mineral soil into the organic soil.

I am finding that [HCO3^-] in the organic soil is on average ~750 mg/L, in comparison with the underlying mineral soil which is around 200 mg/L. I also have evidence that sulfate reduction in the organic layer is occurring, as sulfate concentrations decrease dramatically from mineral soil to the organic soil.

I understand that sulfate reduction produces alkalinity, which should raise the pH. However, the pH in the organic layer, where alkalinity is much higer than the underlying layer, is about 1 pH unit less on average. The average pH in the organic layer is 6.5, and the average pH in the mineral soil is 7.7. So I don't understand why the organic layer has higher pH than the underlying mineral soil if alkalinity is higher. My thinking is that there is some pCO2 increase in the organic soil, but wouldn't that dissolve to form HCO3^-? And consume protons in the process? Thereby increasing the pH.

If anyone can assist me that would be great. Thanks.

[Borek]

To be honest soil chemistry is completely alien to me, so my remarks can be off.

The water is moving from the mineral soil into the organic soil.

So it moves up, not down?

I also have evidence that sulfate reduction in the organic layer is occurring, as sulfate concentrations decrease dramatically from mineral soil to the organic soil.

What are concentrations of sulfate? I mean: you can try to follow the stoichiometry and evaluate amount of H⁺ consumed.

the pH in the organic layer, where alkalinity is much higer than the underlying layer, is about 1 pH unit less on average.

Have you tried to determine ions/substance responsible for the observed alkalinity? Alkalinity is quite a general term, while it is often useful it can be also misguiding.

I don't understand why the organic layer has higher pH than the underlying mineral soil if alkalinity is higher.

That's what I related to  above Alkalinity is related to the buffering capacity. Say, you have two buffer solutions of the same pH, one with a concentration of the buffer twice higher - obviously it will have a higher buffering capacity. It will also have a higher alkalinity, despite having teh same pH. With some juggling and a clever choice of substances involved producing solution that has one unit lower pH but a higher alkalinity/buffer capacity is rather simple.

My thinking is that there is some pCO2 increase in the organic soil, but wouldn't that dissolve to form HCO3^-? And consume protons in the process? Thereby increasing the pH.

No, that's off. CO₂ reacts with water and produces H⁺, not consumes them.

CO₂ + H₂O H⁺ + HCO₃^-

I am afraid that last comment suggests you have a lot to learn about acid-base equilibrium in general, before you can try to analyze the chemistry of your samples

[Loler]

Thanks for your comments. You are right my carbonate chemistry is not adequate but some things:

1) Yes it moves up.
2) I have measurements of sulfate concentrations, so I will do that. 
3) I am finding that both calcium and magnesium are higher in the organic layer. So maybe there is some dissolution of carbonates happening?
4) Not clear what you said.
5) Given that reaction, the H⁺ ion produced when CO₂ dissolves in water is counteracted by the production of bicarbonate? So there is no net generation of H⁺? Is that accurate?

[Borek]

3) I am finding that both calcium and magnesium are higher in the organic layer. So maybe there is some dissolution of carbonates happening?

Doesn't presence of carbonates depend on the soil type? As I wrote earlier, soil chemistry is not my thing, but as far as I am aware the commonly used method of lowering the soil pH is to use some kind of a dolomite granulate, apparently not every soil contains enough of them.

4) Not clear what you said.

I am afraid that's the part you need to learn more about (acid-base equilibrium details, carbonate chemistry is just a part of the more general picture).

5) Given that reaction, the H⁺ ion produced when CO₂ dissolves in water is counteracted by the production of bicarbonate? So there is no net generation of H⁺? Is that accurate?

No, HCO₃^- doesn't counteract H⁺ presence. Pure water has pH of 7.0, open the container and atmospheric CO₂ will dissolve lowering the pH to a bit above 5. That's exactly effect of the reaction I wrote.

[Borek]

5) Given that reaction, the H⁺ ion produced when CO₂ dissolves in water is counteracted by the production of bicarbonate? So there is no net generation of H⁺? Is that accurate?

To add to what I wrote earlier - perhaps you are confused by the fact adding CO₂ changes pH but doesn't change the alkalinity.