Unraveling the Role of Sulfur Compounds in Acid Rain Formation

In the Classroom edited by

Ed Vitz Kutztown University Kutztown, PA 19530

Unraveling the Role of Sulfur Compounds in Acid Rain Formation: Experiments on a Wetted Glass pH Electrode cia H. G. Coelho, and Ivano G. R. Gutz* Fernando S. Lopes, Lu Instituto de Química, Universidade de S~ ao Paulo, Av. Prof. Lineu Prestes, 748. 05508-000, S~ ao Paulo, SP, Brazil *[email protected]

An investigation of the reactions of small molecules provides an important foundation for the evaluation and explanation of the effect of anthropogenic emissions in the troposphere. Daily, thousands of tons of gaseous pollutants and aerosols are emitted to the atmosphere mainly by combustion of fossil fuels by automobiles, trucks, and power plants. Removal of such species from the atmosphere occurs at variable speeds and degrees by physical, chemical, and biological processes followed by dry deposition or precipitation (1). Rainwater analysis has great importance in this context because wet deposition represents an efficient route for the removal of many soluble and particulate atmospheric pollutants. Because of the equilibrium established with the CO2 present in the air, the rainwater in remote unpolluted regions has a pH around 5.6 (2). The acidification of rainwater over many regions of the planet is associated with an increase in the concentration of strong acids (e.g., sulfuric and nitric acids), as well as weak organic acids (e.g., formic and acetic acids) that are emitted directly or formed photochemically. The quantification of sulfur dioxide in the atmosphere represents an important parameter to estimate the contribution of the fossil fuel combustion as well as to evaluate the effectiveness of desulfurization of these fuels. Once emitted to the air, the SO2 can be oxidized in gas or liquid phases, producing the strong acid, H2SO4. In the gas phase, under dry atmospheric conditions, SO2 is oxidized by OH 3 radicals producing the intermediate HSO3 3 and then the product H2SO4. Ozone may play a role in the aqueous-phase oxidation reaction (3), but in slightly acidic media (pH < 4.5), hydrogen peroxide, H2O2, plays the dominant role because its Henry's law constant is much higher than that for ozone (8 104 and 1 10-2 mol L-1 atm-1, respectively). The H2SO4 formed dissociates almost completely, lowering the pH of the liquid droplets (cloud, fog, or rain). Wet deposition is an efficient route for removal of nitrogen emissions including nitric acid and nitrate. However, the oxidation and reduction chemistry of nitrogen oxides occurs predominantly in the gas phase, so it is less suitable for an experiment in which the pH of the liquid phase is used to monitor chemical reactions. This is why it will not be addressed in the demonstration that follows. Nitrogen in a more reduced oxidation state, such as ammonia or amines, is also quite common in the atmosphere. Its effect on rainwater will be demonstrated in this article. If these alkalinizing species are present, the protonation of the base

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occurs to an extent that depends on the concentration and Kb value, with consequent elevation of the pH. Cloud evaporation may thus produce ammonium sulfate aerosol, which is susceptible to dry deposition or action as a condensation nucleus for new clouds (1, 2). Carbonylic compounds such as formaldehyde, acetaldehyde, and acetone, which are commonly present in the troposphere, also participate in several chemical and photochemical reactions in the gaseous and liquid phases. Formaldehyde (CH2O) is emitted during the production and use of household goods (e.g., household hard-surface cleaners and manufactured wood products such as medium-density fiberboard) and combustion processes (e.g., fireplaces and automobile exhaust). CH2O has a lifetime of hours or even days in the air and can oxidize to formic acid or accumulate in the liquid phase owing to its high Henry's law constant (∼6.0 103 mol L-1 atm-1). It can then be removed by precipitation events (1). When CH2O and SO2 are absorbed by clouds, they react in the aqueous phase producing hydroxymethanesulfonic acid HOCH2SO3H (HMSA), a strong, stable, and oxidant-resistant S(IV) species. Acid rain is a popular subject in environmental and general chemistry texts of all levels (4). Laboratory experiments on acid rain formation are available, but many of them present an oversimplified and qualitative view of the involved chemical transformations (typically, detection of the pH change of water by a visual indicator after exposure to combustion gases of sulfurcontaining fuel) or of its impact on limestone or plants (5-10). In this Journal, there are also demonstrations of the acidification of rain by NOx (7, 10). The actual situation is more complex and dynamic and may involve oxidation processes of sulfur or nitrogen oxides in the gaseous and liquid phases, partial neutralization by alkaline species such as ammonia, and inhibition of sulfur oxidation to sulfate by reaction with formaldehyde (11). The occurrence of some of these reactions is demonstrated by measuring with a pH meter the acidity changes of a simulated raindrop during a sequence of experiments.1 To achieve this goal, the wetted hydrophilic surface of the bulb of a glass electrode, universally employed as a potentiometric pH sensor, is used to emulate a raindrop by directly exposing it to the gaseous phase containing the reactants diluted in ambient air, and the pH is recorded as a function of time. A common misconception is that less acidic rainwater is less polluted. However, this is not true, as will be demonstrated by exposing the simulated raindrop on the glass electrode to an

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r 2010 American Chemical Society and Division of Chemical Education, Inc. pubs.acs.org/jchemeduc Vol. 87 No. 2 February 2010 10.1021/ed800046j Published on Web 01/12/2010

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Journal of Chemical Education

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In the Classroom

alkaline compound such as NH3. In urban and rural areas, ammonia emissions can be quite high. Experimental Overview All demonstrations involve potentiometric pH measurements of a water film (with an average thickness of ∼40 μm) on the bulb of a conventional combination glass electrode for pH measurements. Low-cost Quimis electrodes (model QA338-ECV) and a high-quality Metrohm electrode (model 6.0232.100) were tested and found equally suitable. The pH can be measured with any benchtop or portable pH meter or potentiometer. Preferably, but not necessarily, the signal should be recorded with a plotter or digitally acquired with a computer.2 Stoppered 500 mL Erlenmeyer flasks were used to store the solutions and gases used in the demonstration. SO2 was generated by burning a minimum quantity of sulfur. A conical coil was closely wound at one end of long copper wire and filled with ca. 20 mg of sulfur powder. The sulfur powder was ignited with a match and the wire was suspended inside the Erlenmeyer that was stoppered afterward. Gaseous SO2 can also be prepared by reaction of 0.5 mol L-1 NaHSO3 with 3 mol L-1 H2SO4. Three other solutions were prepared and allowed to equilibrate in Erlenmeyer flasks: (a) H2O2 (10% v/v), (b) aqueous NH3 (2% v/v), and (c) CH2O (2% v/v). All flasks except the one for SO2 generation can be prepared hours before the demonstration and kept well stoppered. The experiments can be repeated at will with the same flasks by the lecturer or by students (under supervision). It is important to emphasize that the glass electrode never comes in contact with the reactant solution. All reactions occur in the water layer that wets the glass bulb of the electrode exposed to the gaseous phase in the flasks (Figure 1). Buffer solutions with pH-certified values of 4, 7, and 10 (Merck) were employed to calibrate the glass electrode. A NaCl solution (2 g L-1) was placed in a wash bottle and used instead of pure water to wash the electrode bulb, thus, granting electrolytic conductance to the liquid film on the electrode bulb and on the porous junction of the reference electrode. After each experiment, the glass electrode was rinsed with this electrolyte until the pH reached 5.6 ( 0.2. Immersion of the electrode in a diluted acid solution before rinsing was found to speed up the decontamination of the glass membrane after excursions to high pH. Hazards This experiment is intended as a lecture demonstration as some reagents might be harmful if not handled properly. The experiment must be performed in a laboratory with good ventilation. Sulfur is preferably ignited in a fume hood. Sulfur dioxide vapors are irritating to throat, mucous membranes, and upper respiratory tract. Severe overexposure may result in pulmonary edema, permanent lung injury, or death. Hydrogen peroxide causes eye and skin burns and may cause irritation of the respiratory tract with burning pain in the nose and throat, coughing, wheezing, shortness of breath, and pulmonary edema. Concentrated aqueous ammonia is damaging to eyes. Even contact with dilute aqueous ammonia can lead to serious eye damage. It is toxic if swallowed and is harmful if inhaled or in contact with skin. Formaldehyde is an eye, skin, and respiratory tract irritant. Inhalation of vapors can produce narrowing of the bronchi and an accumulation of fluid in the lungs. In high concentrations, it is a potent sensitizer and a probable human carcinogen. 158


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Vol. 87 No. 2 February 2010

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Figure 1. Combination glass electrode exposed to the SO2(g)-containing flask; flask with H2O2 solution (∼50 mL) in equilibrium with H2O2 in the gas phase; flask with aqueous NH3 in equilibrium with gaseous NH3; and flask with CH2O in equilibrium with the gas phase.

Procedure and Results The electrode is calibrated with standard buffers and rinsed with the NaCl solution used to wet the glass electrode. This solution resembles clouds formed from oceanic air masses, which are characterized by higher NaCl content than continental ones. To speed up the response during exposure to the reactant gases, any drop hanging from the bottom of the glass bulb after rinsing with the electrolyte solution should be removed by gently touching it with absorbent paper so that only a thin liquid film remains on the glass electrode. After exposure to a reactant, fast evaporation of the liquid film in dryer air, causing unreliable pH measurement, can be avoided by keeping the glass electrode in humid air contained in a 2 L beaker with 0.1 L of water.3 Having prepared the Erlenmeyer flasks as described and calibrated the pH sensor and washed it with electrolyte, the demonstration can be conducted. Data acquisition is started by recording the initial pH for 1-2 min before introducing the electrode into the Erlenmeyer with SO2 for 2-3 s. A sharp decrease in the pH value is observed immediately owing to the absorption of SO2. A few seconds after removal of the electrode from the SO2 source, the pH value increases again owing to equilibrium displacement with release of SO2 to the surrounding air, as illustrated in Figure 2, region A, that presents a minimum pH of 3.4 and a recovery with a pH of 4.2. The first dissociation of sulfurous acid is quite extensive in this pH range (H2SO3 presents pK1 = 1.9, pK2 = 7.2), acidifying the medium: SO2 3 H2 O / H2 SO3 / HSO3 - þ Hþ ð1Þ From the pH values, the total dissolved SO2 concentration of all forms in eq 1 can be estimated: Ka1 ¼

½HSO3 - ½Hþ ½Hþ 2 ¼ ½H2 SO3 ½H2 SO3 tot - ½Hþ ½H2 SO3 tot ¼ ½Hþ þ

½Hþ 2 Ka1

ð2Þ

(3Þ

Calculations can be done with a freeware program such as CurTiPot (13) that takes into account that pH is a measure of -log aHþ (and not -log[Hþ]) and generates the distribution

pubs.acs.org/jchemeduc

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r 2010 American Chemical Society and Division of Chemical Education, Inc.

In the Classroom

Figure 2. Screenshot of the experimental results: (A) exposure of the wetted electrode to gaseous SO2 (2-3 s) and recovery in ambient air; (B) rinsing of the electrode bulb with the NaCl solution; (C) exposure to SO2 (2 s) and H2O2 (10 s) and observation in ambient air; and (D) exposure to NH3 (2 s) and follow-up in ambient air.

diagram of the species as well, showing that HSO3- is the predominant species in solution (97% at pH 3.4 and 99% at pH 4.2). At an ionic strength of 0.034 mol L-1 (dictated by the electrolyte, 2 g L-1 NaCl), a γHþ of 0.84 is estimated by the program while the total SO2 concentrations of 4.8 10-4 and 7.5 10-5 mol L-1 satisfy the measured pHs of 3.4 and 4.2, respectively.4 These results indicate that in just 3 min, ∼85% of the SO2 is released from the liquid to the gas phase. After washing the electrode with the electrolyte solution until a pH of 5.4 ( 0.2 is reached (region B of Figure 2), the exposure to SO2 is repeated, and after 2-3 s, the electrode is quickly transferred to the H2O2-containing Erlenmeyer for ∼10 s. This causes an additional decrease of pH to 2.5-3.0, as a consequence of the fast oxidation of S(IV) to S(VI) with formation of a H2SO4, a stronger acid (pK1 ≈ -3, pK2 = 2.0). Another remarkable difference, observed in Figure 2 region C after removal of the H2O2 source, is the stability of the low pH owing to the nonvolatility of the sulfuric acid. The reaction in the liquid phase can be represented as (2): HSO3 - þ H2 O2 ðaqÞ f SO2 OOH - ðaqÞ þ H2 OðlÞ ð4Þ SO2 OOH - ðaqÞ þ Hþ f H2 SO4 ðaqÞ

ð5Þ

The oxidation of SO2 can take place in the gas phase as well, initiated by hydroxyl radicals (OH 3 ) in the presence of a catalyst, M: ð6Þ SO2 þ HO 3 þ M f HOSO2 3 þ M HOSO2 3 þ O2 f SO3 þ HO2 3

(7Þ

SO3 þ H2 O f H2 SO4 ðaqÞ

(8Þ

The effect of NH3 on rainwater is demonstrated by continuing the experiment with an exposure of the glass electrode to the Erlenmeyer containing aqueous NH3 for 2-3 s. The abrupt increase in the pH value followed by stabilization between


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Figure 3. Screenshot of the experimental results: (A) wet electrode exposed to air; (B) exposure to CH2O; (C) observation in ambient air; (D) exposure to SO2 (2 s) and observation for some minutes; and (E) exposure to H2O2 (10 s) and follow-up in ambient air.

8.5 and 9.0 after removal of the NH3 source is apparent in Figure 2, region D. Highly polluted rainwater can, thus, assume pH values in the range 2-11, depending on the reactant in excess. Because of the ionization in aqueous solution, the formation of the crystalline salt (NH4)2SO4 occurs only by evaporation of the water from the droplet, a process that is common in clouds and represented by (2): NH3 3 H2 O / NH4 þ þ OH (9Þ 2NH4 þ þ SO4 2 - f ðNH4 Þ2 SO4 ðsÞ

(10Þ

Conversely, by contacting humid air masses, aerosols such as (NH4)2SO4 or other particulate matter can act as condensation nuclei for water vapor, generating polluted clouds, often at great distances from the sources. The stabilization of S(IV) species by reaction of HSO3with formaldehyde is demonstrated by washing the glass electrode with the NaCl solution and exposing it to the flask containing CH2O vapor for about 10 s before contact with SO2, as shown in step B of Figure 3. Formaldehyde is neither acidic nor basic in water, so that the pH remains practically unchanged during its dissolution in the liquid film on the electrode bulb or afterward (Figure 3, step C). Subsequent exposure to gaseous SO2 for 2-3 s (Figure 3, step D) causes an abrupt decrease in the pH to 3.2 ( 0.2, that remains low even after removing the S(IV) source, resembling the H2SO4 formation. However, no strong oxidant is available and the fast reaction observed is the formation of the nonvolatile sulfonic acid, HMSA, which is a strong acid (pKa ≈ -3). From the SO2 equilibria previously described, HMSA formation can be represented by Hþ þ HSO3 - þ CH2 O f HOCH2 SO3 H f HOCH2 SO3 - þ Hþ

(11Þ

The S(IV) in the HMSA anion is most resistant to oxidation in the pH region of 4-10, but even at pH around 3 (uncommon for acid rain), a high concentration of H2O2 seems to act very slowly on it, as shown in region E of Figure 3.


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In the Classroom

information about this subject. The proposed set of experiments was enthusiastically welcomed by undergraduates as an interesting and meaningful way to understand the underlying chemical reactions and equilibria. In addition to the environmental questions, important topics discussed in undergraduate chemistry courses, such as acid-base and redox equilibria and kinetics, Henry's law, potentiometry, and the Nernst equation, are incorporated in the demonstration and can be explored in class as discussions and exercises. Acknowledgment The authors wish to thank Paulo Alves Porto and coworkers for presenting the demonstration to a class of chemistry undergraduates, to anonymous referees of JCE for reproducing and commenting on the experiments, and to CNPq (Conselho Nacional do Desenvolvimento Científico e Tecnologico, Brazil) and FAPESP (Fundac-~ao de Amparo a Pesquisa do Estado de S~ao Paulo) for a grant and fellowships. Notes Figure 4. Pictorial representation of the combination glass electrode and the most relevant chemical equilibria in a oversized cloud or rainwater drop, involving S(IV) and S(VI) species in the presence of H2O2, NH3, and CH2O. Evaporation of clouds will generate, for example, (NH4)2SO4 aerosol, as denoted with the yellow line. Dissolved NaCl (electrolyte) was omitted for clarity.

Students' understanding can be reinforced with a simple test tube experiment: A solution of HMSA is prepared by mixing a sodium sulfite solution with formaldehyde in excess. Drops of H2O2 are added and followed by addition of calcium or barium nitrate. The solution in the test tube remains clear for a long time. If the experiment is repeated without formaldehyde, the solution turns cloudy from the calcium or barium sulfate precipitate. Whereas the volume of the liquid film on the electrode (estimated by weighing as ∼40 μL) is in the size range of raindrops, the thickness of the film on the glass surface, averaging 40 μm, is comparable to the diameter of cloud and fog droplets (1-100 μm). Because of their higher surface-to-volume ratio and greater residence time, favoring gas-liquid equilibria, fog and clouds are more efficient pollutant scavengers than rainwater drops. To enhance the effects observed in the demonstration, the concentrations of the gaseous reactants were orders of magnitude higher than those found in atmospheric air (mmol per L versus nmol per L, respectively). However, the main reactions, summarized in Figure 4, are the same and the experiments are representative of acid rain formation and neutralization. Finally, it should be emphasized to the students that there are other contributors to the acidification of the wet deposition, such as HNO3 and organic acids, originating from emissions of nitrogen oxides, carbonylic compounds, and carboxylic acids.

Literature Cited

Conclusions The phenomenon of acid deposition, best known as acid rain, is a serious concern in many regions of the planet. Students are widely exposed to controversial, incomplete, or even wrong 160

1. Most of the experiments were performed by undergraduates at the authors' institute or were demonstrated to the participants of the S~ao Paulo State Chemistry Olympiads, a classification step for the Brazilian Chemistry Olympiad, whose national champions participate in the International Chemistry Olympiads (12). 2. This is a simple matter now that most potentiometers present a serial or USB communication port and older ones have an analog output that can be connected to a computer through any data acquisition board (DAQ) with a 12 bits or better analogto-digital converter. The DAQ chosen here was a compact external unit that connects to the computer through a USB port, the USB-6009 from National Instruments. It presents a resolution of 14 bits and a sampling rate of 48 ksamples/s, higher than needed, so that averages of the measurements were stored and plotted at a frequency of 1 Hz. Straightforward software developed in LabView (available from the authors on request) served to convert the potential difference of the electrodes into pH after calibration with buffers, and to store and display the data. Eventually, the authors substituted the commercial pH meter by a single high-impedance operational amplifier (OA model LF356) wired as a voltage follower. The noninverting input of the OA was connected to the glass electrode; the inverted input, to the output pin of the OA and to an input of the DAQ; and the reference electrode to ground. The pH meter is also superfluous in laboratories having a DAQ with high impedance analog input such as the Dataquest from Vernier or SensorDAQ, both compatible with LabView. 3. It is assumed that the relative humidity in the laboratory is above 50%. 4. By assuming pH = -log[Hþ], 4.1 10-4 and 6.3 10-5 mol L-1 would be obtained.


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1. Mozeto, A. A. Quím. Nova Escola 2001, 41–49. 2. Seinfield, J. H.; Pandis, S. N. Atmospheric Chemistry and Physics; From Air Pollution to Climate Change; John Wiley & Sons: New York, 2006; pp 306-308.


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In the Classroom

3. Bunce, N. J. Introduction to Environmental Chemistry; Wuerz Publishing Ltd.: Winnipeg, Canada, 1993. 4. Cooper, M. M.; Elzerman, A. W.; Lee, C. M. J. Chem. Educ. 2001, 78, 1169–1169. 5. Powers, D. C.; Higgs, A. T.; Obley, M. L.; Leber, P. A.; Hess, K. R.; Yoder, C. H. J. Chem. Educ. 2005, 82, 274–277. 6. Goss, L. M.; Eddleton, J. E. J. Chem. Educ. 2003, 80, 39–40. 7. Driscoll, J. A.; Jones, R. F. J. Chem. Educ. 1997, 74, 1424– 1425.


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Ophardt, C. E. J. Chem. Educ. 1985, 62, 257–257. Zajicek, O. T. J. Chem. Educ. 1985, 62, 158–159. Gleason, G. I. J. Chem. Educ. 1973, 50, 718–719. Rocha, F. R.; Coelho, L. H. G.; Lopes, M. L. A; Carvalho, L. R. F.; Fracassi da Silva, J. A.; do Lago, C. L.; Gutz, I. G. R. Talanta 2008, 76, 271–275. 12. Allchemy Home Page. http://allchemy.iq.usp.br (accessed Jul 2008). 13. CurTiPot Freeware Page. http://www2.iq.usp.br/docente/gutz/ Curtipot_ (accessed Mar 2009).

8. 9. 10. 11.


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Unraveling the Role of Sulfur Compounds in Acid Rain Formation

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