The Impact of pH Change on Acid Deposition, Infrastructure, and Water Sources

Acid Deposition Inquiry

Background Information

In recent years, acid deposition has received more and more attention as its harm to the environment becomes obvious. Acid deposition is defined as the accumulation of acids or acidic compounds on the surface of the Earth. It can occur in bodies of water, infrastructure, or plant life. Acid deposition often harms the environment: it can kill fish by altering the pH of lakes, weaken trees, etc. Wet acid deposition refers to acidic precipitation, which comes in the form of rain, snow, or fog.

Wet acid deposition can also harm man-made buildings. It can corrode metal and destroy paint and stone. A noteworthy example is India’s Taj Mahal, which is made of marble. Its white color has slowly deteriorated due to acid rain, and the monument now appears pink.

Acids are defined as proton donors in Brønsted-Lowry theory. However, in order for a proton to be donated, then there must be another substance available to accept the proton. This substance is known as a base.

Consider the generic acid-base reaction between an acid HA and base B:

HA + B ↔ A- + BH+

HA is the acid here, as it donates a proton to B, the base which accepts the proton.

The pH scale is a quantitative scale of acid strength. It is based on the concentration of hydrogen ions, specifically pH = -log10 [H+]. As the pH decreases, the more acidic a substance is.

Even a minor change in pH can prove deadly. For example, the pH of human blood is 7.4, and this is strictly maintained. Even a slight disruption will be fatal. Acid rain usually takes on a pH of 5.5 or so, which may not seem significantly different from the pH of normal water but which can severely damage the world we live in.

Acid rain is mostly acidic due to the H2SO4 and HNO3 it contains. Pollution in the form of SO3 and NO2 goes up to the atmosphere, where they react with water in the following reactions:

SO3 + H2O → H2SO4

NO2 + OH- → HNO3

Acid deposition contains seemingly miniscule amounts of sulfuric and nitric acid, but the implications can be huge for environment and infrastructure.

Purpose

The purpose of this lab is to investigate the effect of acid deposition on infrastructure as well as various sources of water.

Research Question

How does acid deposition affect other sources of water as well as infrastructure in terms of pH change?

Variables

The independent variables in this lab are the concentration of H2SO4 as well as the three different water sources. The concentrations of H2SO4 will be 1.00 M, 0.75 M, 0.50 M, 0.20 M, 0.10 M, and 0.05 M. They will be made from a stock solution of 1.00 M H2SO4. Since M = moles/liter, the stock solution will be diluted properly using distilled water. The stock solution will be kept as is for 1.00 M. For 0.75 M, 150 mL of stock solution will be measured out and 50 mL of distilled water will be added. For 0.50 M, 100 mL of stock solution will be measured out and 100 mL of distilled water will be added. For 0.20 M, 40 mL of stock solution will be measured out and 160 mL of distilled water will be added. For 0.10 M, 20 mL of stock solution will be measured out and 180 mL of distilled water will be added. For 0.05 M, 10 mL of stock solution will be measured out and 190 mL of distilled water will be added. These measurements will be carried out using a beaker and graduated cylinder.

The three different water sources are rainwater from Portland, OR; tap water from the Tualatin Water Valley District; and creek water from Mill Pond. Approximately a liter of each source will be collected. The rainwater will be collected using a pan to catch rainfall, the tap water will be collected from the sink, and the creek water will be collected directly from the creek.

The dependent variable in this lab is the pH. The pH of the different concentrations of H2SO4 as well as the pH of the three different water sources will be measured. In addition, the pH of the different independent variables when mixed with CaCO3 will be measured. It will be measured using a Vernier pH Sensor.

The constants in this lab include the pH sensor used, the amount of CaCO3 used for each variable, the substance (CaCO3) mixed with each variable, the acid (H2SO4) used, and all glassware.

Materials

• 1 Vernier pH Sensor
• 200 mL each of 1.0 M, 0.75 M, 0.50 M, 0.20 M, 0.10 M, and 0.05 M H2SO4
• 1 L rainwater (Portland, OR)
• 1 L tap water (Tualatin Valley Water District)
• 1 L creek water (Mill Pond)
• Approx. 200 g CaCO3 (marble chips)
• Mortar and pestle
• 50 mL graduated cylinder
• 250 mL beaker
• 50 mL beaker
• 680 mL distilled water
• Electronic scale

Procedure

1. Gather all materials.
2. Using the mortar and pestle, crush up 400 g of CaCO3.
3. Measure the pH of 200 mL of 1.0 M H2SO4 using the Vernier pH Sensor. Record.
4. Using the electronic scale, measure out 10 g of CaCO3.
5. Using the graduated cylinder, measure out 50 mL of 1.0 M H2SO4 and place in 50 ml beaker.
6. Place 10 g of CaCO3 in the same 50 mL beaker.
7. Measure pH of the substances inside the 50 mL beaker. Record.
8. Repeat steps 3 - 7 with 0.75 M, 0.50 M, 0.20 M, 0.10 M, and 0.05 M H2SO4.
9. Repeat steps 3 - 7 with rainwater (Portland, OR), tap water (Tualatin Valley Water District), and creek water (Mill Pond).
10. Repeat steps 3 - 9 twice.

Raw Data Tables

Trial 1

Concentration of H2SO4 (M) pH of H2SO4 solution

+/- 0.01 pH of H2SO4 solution + CaCO3

+/- 0.01

1.0 0.62 1.13

0.75 0.75 1.36

0.50 0.82 2.45

0.20 0.95 1.55

0.10 1.03 4.97

0.05 1.17 5.38

Trial 2

Concentration of H2SO4 (M) pH of H2SO4 solution

+/- 0.01 pH of H2SO4 solution + CaCO3

+/- 0.01

1.0 0.64 1.29

0.75 0.80 1.83

0.50 0.84 2.42

0.20 1.01 2.05

0.10 1.03 5.27

0.05 1.32 6.12

Trial 3

Concentration of H2SO4 (M) pH of H2SO4 solution

+/- 0.01 pH of H2SO4 solution + CaCO3

+/- 0.01

1.0 0.67 1.61

0.75 0.72 1.70

0.50 0.79 2.45

0.20 0.99 2.23

0.10 1.03 4.99

0.05 1.31 5.04

Trial 1

Water source pH of water

+/- 0.01 pH of water + CaCO3

+/- 0.01

Rainwater (Portland, OR) 5.80 6.13

Tap water (Tualatin Valley Water District) 7.48 8.01

Creek water (Mill Pond) 6.52 8.46

Trial 2

Water source pH of water

+/- 0.01 pH of water + CaCO3

+/- 0.01

Rainwater (Portland, OR) 5.82 6.44

Tap water (Tualatin Valley Water District) 7.30 8.95

Creek water (Mill Pond) 5.97 7.21

Trial 3

Water source pH of water

+/- 0.01 pH of water + CaCO3

+/- 0.01

Rainwater (Portland, OR) 5.56 7.03

Tap water (Tualatin Valley Water District) 6.68 8.47

Creek water (Mill Pond) 5.90 7.11

Analysis

Averages:

Independent Variable pH +/- 0.01 pH when CaCO3 is added +/-0.01

1.0 M H2SO4 0.64 1.34

0.75 M H2SO4 0.76 1.63

0.50 M H2SO4 0.82 2.44

0.20 M H2SO4 0.98 1.94

0.10 M H2SO4 1.03 5.08

0.05 M H2SO4 1.27 5.51

Rainwater 5.73 6.53

Tap water 7.15 8.48

Creek water 6.13 7.59

Example calculation:

Average for 1.0 M H2SO4

(0.62 + 0.64 + 0.67)/3 = 0.64

Average for 1.0 M H2SO4 with CaCO3

(1.13 + 1.29 + 1.61)/3 = 1.34

Concentration of H3O+ released

Independent Variable Concentration of H3O+

+/- 0.01 Concentration of H3O+ when CaCO3 is added

+/-0.01 Difference in concentration of H3O+ when CaCO3 is added

+/- 0.01

1.0 M H2SO4 2.29 * 10-1 4.57 * 10-2 2.00 * 10-1

0.75 M H2SO4 1.74 * 10-1 2.34 * 10-2 1.35 * 10-1

0.50 M H2SO4 1.51 * 10-1 3.63 * 10-3 2.40 * 10-2

0.20 M H2SO4 1.05 * 10-1 1.15 * 10-2 1.10 * 10-1

0.10 M H2SO4 9.33 * 10-2 8.32 * 10-6 8.91 * 10-5

0.05 M H2SO4 5.37 * 10-2 3.09 * 10-6 5.75 * 10-5

Rainwater 1.86 * 10-6 2.95 * 10-7 1.58 * 10-1

Tap water 7.08 * 10-8 3.31 * 10-9 4.68 * 10-2

Creek water 7.41 * 10-7 2.57 * 10-8 3.47 * 10-2

Example calculations:

Concentration of H3O+ for 1.0 M H2SO4:

pH = - log10 [H3O+]

0.64 = - log10 [H3O+]

[H3O+] = 10-0.64 = 2.29 * 10-1

Concentration of H3O+ for 1.0 M H2SO4 with CaCO3:

pH = - log10 [H3O+]

1.34 = - log10 [H3O+]

[H3O+] = 10-1.34 = 4.57 * 10-2

Multiplied difference in concentration of H3O+ for 1.0 M H2SO4 compared to concentration with CaCO3:

1.34 - 0.64 = 0.70 = - log10 [H3O+with CaCO3 /H3O+without CaCO3]

[H3O+] = 10-0.70 = 2.00 * 10-1

Graphs

Conclusion

In this inquiry, the pH of different concentrations of H2SO4 was measured, as well as the pH of different sources of water. Theoretically speaking, the pH of an acid should not change as it is diluted with distilled water, since the pH measures concentration of H3O+ ions released. However, there was a clear trend in pH increasing as the H2SO4 was diluted. This may be due to the fact that the solutions were not mixed thoroughly, and thus the dispersion of H3O+ ions was not equal throughout the beaker the solution was in. This can be alleviated next time by using a stirring rod and waiting a bit longer before measuring the pH with the pH sensor.

In addition, the concentrations of H2SO4 used were much too high to simulate acid rain accurately. Acid rain has very small amounts of sulfuric acid--nothing close to 1.0 M. It may be more accurate to use concentrations of H2SO4 such as 0.001 M.

Also, the addition of CaCO3 to H2SO4 solutions was ten grams to each, regardless of the number of moles of H2SO4 in the solution. If this experiment were to be carried out again, it would be better to add the same number of moles of CaCO3 as H2SO4, so there are no limiting reagents. CaCO3 and H2SO4 react together in the following reaction:

H2SO4 + CaCO3 → CaSO4 + H2O + CO2

In this inquiry, it was seen that the effects of acid deposition can be sobering. The pH of CaCO3 in water is generally about 9. However, once mixed with acidic substances, the pH drops quickly. In reality, this translates to the erosion of buildings made of marble or limestone, although at a decidedly less dramatic rate than seen here. Acid rain, as measured here, has a pH of about 5.7. While this may not seem much different from distilled water at 7, over years it can destroy infrastructure as well as environments.

It isn’t possible to know exactly the effect of acid deposition on the environment simply from this inquiry--while the pH of creek water was measured, there isn’t a way to compare it to the optimum value, as creeks and ponds hold life. The respiration and photosynthesis of the life inside these bodies of water contribute to the pH, and thus it cannot be claimed simply that 7 is the optimum pH for a creek. An experiment may be carried out where the pH of the creek is continuously monitored over a period of years, to study the effect of acid deposition on the environment.

Unsurprisingly, the pH of a substance versus the pH once CaCO3 is added follows a logarithmic trend. In addition, the concentration of H3O+ ions released by a substance versus the concentration of H3O+ ions released once CaCO3 is added follows an exponential trend. Since pH is calculated using logarithms, this is expected. However, it also means that the slightest shift in pH translates to a larger difference in number of H3O+ ions being released, so it is imperative that pH stays stable in environments.

Possible sources of error in this experiment include human error, systematic error of the pH sensor, and fluctuations in temperature. CaCO3 was ground up using a mortar and pestle, which may have allowed some larger pieces of marble chips to not be crushed up completely and thus were more difficult to dissolve in solution. And as always, more trials would improve the accuracy of this inquiry. In future experiments, it may be of note to study possible buffers to counteract the effects of acid deposition on the environment, or the effects of acid deposition on other building materials besides CaCO3.