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Carbon Species of the Lithosphere and Hydrosphere: the Carbonic Acid System1

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Carbon Species of the Lithosphere and Hydrosphere: the Carbonic Acid System1 Experiment 18B SC151 10/07/10 CARBON SPECIES OF THE LITHOSPHERE AND HYDROSPHERE: 1 THE CARBONIC ACID SYSTEM MATERIALS: pH meter (2); magnetic stirrer and stirbar (2); standardization buffers (pH=4, 7 and 10); 50 mL buret (2); 400 mL beaker (3); 100 mL beaker (2); 100 mL graduated cylinder; filter funnel and filter paper; anhydrous Na2CO3; saturated solution of MgCO3; 0.036 M H2SO4; 0.0072 M H2SO4. PURPOSE: The purpose of this experiment is to introduce the use of the pH meter and to identify important features in the titration curves of weak acids and bases. LEARNING OBJECTIVES: By the end of this experiment, the student should be able to demonstrate the following proficiencies: 1. Standardize and use a pH meter. 2. Perform a pH titration. 3. Identify the dominant chemical species at various points in a titration curve. 4. Obtain values of pKa for weak acids from pH titration data. DISCUSSION Every carbon atom in your body has been used in countless other molecules since the beginning of time, as CO2 in the air you breathe, as carbohydrates in the food you eat, as bicarbonate “fixed” by sea life, or as carbonate salts in the rocks and soil you walk on. The Carbon Cycle is a complex series of processes which describes how carbon atoms circulate among these realms of atmosphere, biosphere, hydrosphere and lithosphere. By far, the greatest store of carbon atoms on earth can be found in marine sediments and sedimentary rocks (~8 x 1016 metric tons), with the ocean providing a distant second store (~4x1013 metric tons). Only a relatively small amount is found in the atmosphere (~7x1011 metric tons)2. The carbon in these three realms is connected through the equilibria of carbonic acid and its salts. Carbonic acid, H2CO3, in the hydrosphere is formed when atmospheric carbon dioxide dissolves in sea water: CO2 (g) + H2O (l) ⇌ H2CO3 (aq) Dissolved carbonic acid is a typical weak polyprotic acid, the principal equilibria being: + 7 H2CO3 (aq) + H2O (l) ⇌ H3O (aq) + HCO3¯ (aq) Ka1 = 4.2 x 10¯ + 2 11 HCO3¯ (aq) + H2O (l) ⇌ H3O (aq) + CO3 ¯ (aq) Ka2 = 4.8 x 10¯ The dominant aqueous carbon species will depend on the pH of the surrounding waters. At low pH, the carbonic acid form is dominant, while at high pH, carbonate ion is the most important species. The carbonate ion also provides the connection between hydrosphere and lithosphere, in the form of insoluble CaCO3 shells formed by many marine organisms and insoluble MgCO3 formed from sea salts. The carbonic acid system plays an important role in the chemistry of natural waters, both as a “sink” for carbon atoms and as a buffer against acid assaults from acid rain, mine run-off, etc. (Interestingly, the carbonic acid system is one of the most important buffers in blood as well.) In this experiment, we will map out the composition of the carbonic acid system as a function of pH by performing a pH titration. This will allow us to determine the dominant carbon species at any pH, and identify important features in the behavior of polyprotic acids in general. Because it is experimentally much easier to start with a known, fixed concentration of carbonate (or bicarbonate) salt, rather than carbonic acid, we will titrate these basic salts with standard acid solution. Since sulfuric acid, H2SO4, is the dominant acid in acid rain, that will be used as the titrant. The results of this experiment will thus be directly applicable to understanding the function of buffers in soils and natural waters. 1adapted from D.C. Powers, et al, J. Chem. Educ. 82, 274 (2005). 2 M. Pidwirney, Encyclopedia of Earth, http://www.eoearth.org/article/Carbon_cycle; accessed 25 June 2007. E18B-1 PROCEDURE (work in pairs): Part A. Standardization of the pH meter. 1. Your Instructor will demonstrate the use of the pH meter. Pay close attention to the procedure, especially the need to treat the electrodes gently, wash them with water and gently pat dry prior to each measurement, and to avoid leaving the electrodes exposed to air for long periods. Record notes as necessary, for you will use this instrument again. 2. Standardize the pH meter with the three buffer solutions provided (pH = 4, 7 and 10). It is NOT necessary to pour the buffers out of their vials; simply make the measurements in the original containers. Part B. Titration of Carbonate Solutions. Each midshipman will do one of the two titrations listed below, so the pair of lab partners will generate all of the data for both solutions. Note that both experiments will use sulfuric acid as titrant, but at different concentrations. Be sure to use the solution appropriate for your titration. Solution Titrant Na2CO3 0.036 M H2SO4 MgCO3 0.0072 M H2SO4 Titration of Na2CO3 Solution: 1. Weigh out 0.15 g of anhydrous Na2CO3. Place the solid sample in a clean 400 mL beaker, and add 200 mL of distilled water. Add a stir bar, place the beaker on the magnetic stirrer, and stir until fully dissolved. Place the pH electrode in the solution and continue to stir gently. 2. Rinse the buret with 0.036 M H2SO4 solution, remove any bubbles from beneath the stopcock, and refill the buret. Mount the buret in a buret clamp and arrange the buret so that additions of titrant can be made while the pH electrode remains in the solution in the beaker. 3. Record the pH of the solution before any titrant is added. Then, add titrant in 1.00 mL increments until 45 mL of titrant have been delivered. Wait until the pH reading stabilizes before recording the value, or adding more titrant. (This might take a few seconds, or a few minutes, depending on where you are in the titration.) 4. After you have added 45 mL of titrant, remove, rinse and secure the pH electrode, retrieve the stir bar and wash all of your glassware. Excess solutions can be safely discarded in the sinks. Titration of MgCO3 Solution: 1. Place about 200 mL of saturated MgCO3 solution into a clean beaker. To ensure that no solids were included, set up a filter funnel and filter this solution into another clean 400 mL beaker. Add a stir bar and place the beaker on the magnetic stirrer. Place the pH electrode in the solution and begin to stir the solution gently. 2. Rinse the buret with 0.0072 M H2SO4 solution, remove any bubbles from beneath the stopcock, and refill the buret. Mount the buret in a buret clamp and arrange the buret so that additions of titrant can be made while the pH electrode remains in solution. 3. Record the pH of the solution before any titrant is added. Then, add titrant in 1.00 mL increments until 45 mL of titrant have been delivered. Wait until the pH reading stabilizes before recording the value, or adding more titrant. (This might take a few seconds, or a few minutes, depending on where you are in the titration.) 4. After you have added 45 mL of titrant, remove, rinse and secure the pH electrode, retrieve the stir bar and wash all of your glassware. Excess solutions can be safely discarded in the sinks. E18B-2 DATA SHEET – EXP 18B Titration of Na2CO3 with 0.036 M H2SO4 Titration of MgCO3 with 0.072 M H2SO4 Vol. acid Vol. acid Vol. acid Vol. acid pH pH pH pH (mL) (mL) (mL) (mL) E18B-3 Part C. Obtaining Data for Other Carbonate Solutions. 1. From your lab partner, obtain titration data for the solution you did not work with. Each partner will need all of the data for both Na2CO3 and MgCO3 solutions to perform the analysis. 2. Titration data for NaHCO3 solution is provided below. Acid volumes are provided in mL. You will need these data for the analysis. Titration of 0.0071 M NaHCO3 solution with 0.036 M H2SO4 Vol acid pH Vol acid pH Vol acid pH Vol acid pH Vol acid pH 0.0 8.64 11.00 6.39 21.00 3.54 31.00 2.40 41.00 2.13 1.00 7.85 12.00 6.30 22.00 3.20 32.0 2.36 42.00 2.11 2.00 7.52 13.00 6.22 23.00 2.99 33.00 2.33 43.00 2.09 3.00 7.27 14.00 6.11 24.00 2.86 34.00 2.30 44.00 2.08 4.00 7.08 15.00 5.99 25.00 2.75 35.00 2.27 45.00 2.06 5.00 6.95 16.00 5.86 26.00 2.67 36.00 2.24 6.00 6.84 17.00 5.72 27.00 2.60 37.00 2.22 7.00 6.74 18.00 5.51 28.00 2.54 38.00 2.19 8.00 6.68 19.00 5.21 29.00 2.49 39.00 2.17 9.00 6.59 20.00 4.49 30.00 2.44 40.00 2.15 10.00 6.53 DATA ANALYSIS: 1. Use Excel to create separate plots of pH vs. volume H2SO4 for each of the three titrations: Na2CO3 solution, MgCO3 solution and NaHCO3 solution (data provided above). Print these out as full page plots. 2. On each graph, identify the following: i. the equivalence points (there will be two equivalence points in the Na2CO3 and MgCO3 titrations) 2 ii. the form (or forms) of carbon (i.e., CO3 ¯, HCO3¯ or H2CO3) at: a.
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