Erosion of Carbonate Stone by Acid Rain

We have all seen pictures of the devastation to national and international monuments attributed to acid rain. In this experiment we will examine the effect pH has on the amount of loss of material from various types of stone containing calcium carbonate.

Rain water in equilibrium with atmospheric CO₂ at STP has a pH of 5.6. Natural sources of acidic gases (e.g. volcanic emissions of SO₂ and organic acids may serve to reduce local pH values to between 5.0 and 5.6. In acid rain the primary contributions of hydrogen ion besides the natural sources of acidity are sulfurous, sulfuric, and nitric acids, which lower the pH of rain and accelerate weathering processes. The average pH of most rainfall in the eastern United States is between 3.9 to 4.5.

Under equilibrium conditions, the incremental impact of hydrogen ion in pH 4.0 rainwater is expected to be small because the equilibrium solubility of calcite is dominated by its reaction with carbonic acid derived from atmospheric CO₂ that has dissolved in rainwater.

CaCO₃ + H₂CO₃* à Ca²⁺ + 2HCO₃^-

This holds even though most (greater than 99% ) of the atmospheric CO₂ that dissolves remains in the form of dissolved gaseous CO₂.

H₂CO₃* = CO₂ (aq) = CO₂ + H₂CO₃

Procedure:

Each team will be assigned a rock material.  Weigh out approximately 3 grams (measure exactly what you have) for each bottle.  Each group will need approximately 30 grams total.

The filters you will be measuring the total solids on need to be dried prior to being weighed.  Stick them in the oven for an hour and then weigh them.

Next, prepare 4 solutions, 1000mL each ranging from pH = 4.8 to 3.0 in increments of approximately 0.5 pH units. Acidify with H₂SO₄. Record the exact pH. Divide each solution in half.  Place the 3 grams of rock material in each of the eight solutions plus 2 500mL bottles of distilled water.  Let the rocks sit in the solutions for a week, stirring occassionally.

The calcium carbonate concentration will be measured with an EDTA titration. (Read pages 314-322 in Harvey.)  Make up a 0.005M solution of EDTA.  You will need to use the sodium salt of EDTA.  You will need approximately 2 liters for the whole experiment.  A solution of standardized EDTA will be titrated against your samples.  To standardize the EDTA, make up a calibration standard of CaCO₃ .  The exact concentration should be known and be in the range of 5 to 40mg/L. The CaCO3 standard should be made up in the following manner.   Weigh 1.000g anhydrous CaCO3 powder into a 500mL erlenmeyer flask.  Place a funnel in the flask neck and add, a little at a time, 6M HCl until all CaCO3 has dissolved.  Add 200mL distilled water and boil for a few minutes to expel all CO2.   Cool, add a few drops of methyl red indicator, and adjust to the intermediate orange color by adding 3N NH4OH or 6M HCl as required.  Transfer quantitatively and dilute to 1000mL with distilled water. 1mL = 1.00mg CaCO3.

Once the standard CaCO3 solution is made, the EDTA can be standardized.  All of your titrations will proceed in the same basic manner.  CaCO3 solution (either sample or standard) is combined with 3mL of NH3/NH4+ buffer and a few drops of calmagite indicator.  EDTA out of the buret is added until the color goes from red to blue.  The volume of EDTA required is noted and the CaCO3 is calculated.

While titrating your samples, the volume of sample used may vary from rock material to rock material (and pH to pH).  Adjust your volume accordingly.  Ideally the volume of sample that is used should require between 5 and 25mL of EDTA to titrate.

Before you start titrating your samples and after the rocks have sat in the solutions for a week, you will need to filter the samples.  Filter them through the pre-weighed and pre-dried filters.  Dry and weigh the filters after they have been used also.  Divide the filtered solutions into thirds and titrate.  You should end up with a total of 30 solutions that need to be titrated.

Write-up

1) Plot the hydrogen ion concentration versus carbonate ion concentration. What is the slope of your line? What is the significance of this slope? Compare your results with your material versus the material other people used in the class and make a graph with everyone’s data. Was there any statistically significant difference? Explain that difference.

2) Make a similar plot of total suspended solids versus hydrogen ion concentration. Comment on these results.

3) Clean water has a pH of 5.6. Prove this using the following equilibria:

CO₂(g) à CO₂(aq) K_H = 10^-1.5 and

The acid dissociation constant for carbonic acid found in Appendix G,

The partial pressure of carbon dioxide in the atmosphere is 10^-3.5.

4. What do the results indicate? What type of material will be least susceptible to acid rain? Most? If possible, measure the pH of precipitation here in Winona (this needs to be done as soon as possible after the precipitation is collected. What does this indicate in terms of this lab. Summarize the findings of this experiment in terms of damage acid rain can do.

Pre-lab assignment

This is a three week lab, determine the best way to divide up your time between the three weeks. Write it out in your notebook.

How many grams of ethylene diamine tetraacetic acid disodium salt do you need to make 2 liters of a 0.005M solution?

What is the purpose of adding the NH₃/NH₄⁺ buffer to the EDTA titration?

For further information about acid rain and its effects see Acid Rain Page