In the Laboratory
Gasometric Determination of CO2 Released from Carbonate Materials Johan Fagerlund* and Ron Zevenhoven Thermal and Flow Engineering Laboratory, Åbo Akademi University, Turku, Finland *
[email protected] € ran Hulde n and Berndt So € dergård Stig-Go Combustion and Materials Chemistry, Åbo Akademi Process Chemistry Centre, Turku, Finland
As part of the ongoing mineral carbonation research, an inexpensive and simple method for determining the carbonation content of various carbonated materials (mainly carbonated brucite, the mineral form of magnesium hydroxide) was required. Mineral carbonation is part of a larger carbon dioxide capture and storage (CCS) portfolio and has received comparatively little attention despite its inherent potential (natural process, abundant resources, and permanent CO2 storage). For more information about this area, see, for example, the latest literature reviews (1, 2). The method described here is of value to mineral carbonation research owing to its simplicity, low cost, and accuracy, but it could also serve another purpose; that is, to expose undergraduate students to a hands-on implementation of the ideal gas law and some basic inorganic chemistry. In addition, we consider aspects that could easily be overlooked in a theoretical approach, such as the influence of the heat generated during the reaction between a solid carbonate and hydrochloric acid. In a classroom, these aspects could be considered beforehand, for example, in groups, but it could also be useful to allow the students to find these parameters by simple trial and error. For instance, one student might take the sample container into his or her hand without realizing how much it affects the temperature inside the container before he or she compares the obtained result with that of others. The method presented here is based on the reaction between a carbonated material and an aqueous solution of HCl, which generates CO2. This method is not new and literature dating back to the late 1940s has described this method (3). The reaction system is sealed and the pressure increase is used as a measure of the carbonation degree of the sample, as has also been noted by others more recently (4-6). Choi and Wong used a similar method to determine the reaction kinetics and carbonate content of eggshells, but did not assess the accuracy of the method (4). Roser and McCluskey demonstrated the use of pressure measurement and data logging to determine the stoichiometry of reactions between various carbonates and HCl (5). Horvath et al. was able to accurately determine the carbonate content of soils measuring the pressure increase inside a sealed vessel (6). The method described here takes a slightly different approach, combining several aspects not described in any single literature source cited above. Pile et al. from the Los Alamos National Laboratory (LANL) developed a method for determining the carbonate content of mineral carbonates (7). They measure the carbonate content indirectly through volume change and hydrostatic 1372
Journal of Chemical Education
_
_
pressure; however, we measure the pressure directly. Despite the simple design and reported high accuracy of the LANL system, we were unable to accurately repeat the tests on materials that may be low in carbonate and high in magnesium hydroxide and decided to use (similarly to refs 4-6)) the method described here. Design and Construction of Apparatus The equipment used to determine the CO2 release of a material is simple and available in most teaching laboratories, the most important component being a pressure gauge. Any pressure measurement device capable of determining the pressure to a precision of about 100 Pa could be used. In principle, a manometer could suffice, although we used a PASPORT Absolute Pressure/ Temperature Sensor, PS-2146. Together with PASCO's software called DataStudio, the pressure (and temperature) can be logged and displayed simultaneously as the experiment proceeds. This equipment is especially developed for educational purposes. In addition to a pressure gauge, a container for the sample, a smaller container for the acid, an airtight plug, and some flexible tubes to attach to the pressure gauge are needed. The container for the sample needs to be durable and transparent, such as a sturdy glass jar as the pressure builds up inside. An illustration of the experimental setup is shown in Figure 1. A magnetic stirrer is not shown as it is not necessary; however, it is preferable for more consistent results. The system should not be changed after it has been assembled as any change requires the system to be recalibrated (see the calibration section). The temperature sensor should be located inside the glass jar. It is also possible to assume constant temperature, but this gives less accurate results. Experimental Procedure The carbonated sample (0.1-0.5 g) is accurately weighed ((0.1 mg) using a precision scale. About 2 mL of 5 M HCl (overstoichiometric) is placed in a small cup that fits inside the glass jar. The carbonate sample and the HCl container are carefully inserted into the glass jar together with a magnetic stir bar. The glass jar is sealed with a rubber cap containing the temperature sensor and a tube that is connected to a pressure gauge. The data logger (if available) is started and the system is stabilized (humidity builds up) for a short period of time. The magnetic stirring device is started, which spills the HCl solution, and the reaction is left to go to completion, that is, until no more pressure increase is observed.
_
Vol. 87 No. 12 December 2010 pubs.acs.org/jchemeduc r 2010 American Chemical Society and Division of Chemical Education, Inc. 10.1021/ed1001653 Published on Web 10/08/2010
In the Laboratory
Figure 1. Experimental setup.
The experiment takes less than 10 min, but it is important that the system is allowed sufficient time to reach thermodynamic and chemical equilibrium before and after the stirring, ensuring that all carbonates have dissolved. The steady state of the system can be determined using the pressure gauge. If the pressure is stable, so is the system. It is helpful for the students to see the progress of the reaction on a computer screen (examples in Figure 4). Hazards Hydrochloric acid is corrosive and causes burns to all body tissue. Calculations Using the ideal gas law with a constant volume system, one of remaining three variables can be determined knowing the other two variables. In this experiment, the temperature and the pressure are measured so the only unknown variable is the amount of the substance. Assuming that the created atmosphere consists mostly of CO2 and a small quantity of water, the mass of CO2 released from the sample can be calculated. It is also worth mentioning that the dissolution of CO2 into the solution is negligible owing to the high acidity of the solution and the magnetic stirring. The reaction taking place inside the reaction vessel is MgCO3 ðsÞ þ 2HClðaqÞ f Mg 2þ ðaqÞ þ 2Cl - ðaqÞ þ H2 OðaqÞ þ CO2 ðgÞ
ð1Þ
Similarly, knowing that most of the samples also contain magnesium hydroxide or magnesium oxide, the reaction equations for these with HCl are MgOðsÞ þ 2HClðaqÞ f Mg2þ ðaqÞ þ 2Cl - ðaqÞ þ H2 OðaqÞ ð2Þ MgðOHÞ2 ðsÞ þ 2HClðaqÞ f Mg 2þ ðaqÞ þ 2Cl - ðaqÞ ð3Þ
þ 2H2 OðaqÞ
Figure 2. Relative humidity as a function of time inside the glass jar for two different experiments. Note the different time scales of the two images.
increase the system pressure to a small degree owing to the inevitable vaporization of some of the water (HCl vaporization contributes only a very small fraction of the water partial pressure and is therefore neglected, see the supporting information and ref 8). The vapor content increases can be seen as a small, yet steady, increase in the pressure after the glass jar has been sealed. And it can be shown (Figure 2) that most of the water evaporating from the HCl solution takes place before the reaction (typically around 500 s) starts. Allowing the system to stabilize before the reaction eliminates the small equilibrium increase in pressure from H2O. However, depending on the composition of the sample (e.g., MgO, Mg(OH)2, or MgCO3) the system will behave differently. The dissolution of MgO and Mg(OH)2 is more exothermic (9.3 and 6.7, respectively) than that of MgCO3; thus, more heat is released when either MgO or Mg(OH)2 is in the sample. This effect can clearly be seen from the data in Figure 2. In the case of the sample containing only MgCO3, the relative humidity increase due to the dissolution reaction is barely visible (around 1000 s), whereas in the case of a mixture of MgCO3 and Mg(OH)2, the increase in both temperature and relative humidity is obvious (around 1600 s). In other words, water vapor plays a less significant role in samples with a high carbonation degree, whereas samples with a lower carbonation degree can easily be overestimated. The relative humidity was measured by inserting a measuring (temperature, humidity) probe into a glass jar, similar to an actual experiment run. The carbonation degree is determined by comparing the amount of CO2 released from the sample to the calculated release from a 100% pure carbonate of the same species and mass. The amount of CO2 in a 100% pure carbonate sample is calculated knowing the molar mass of the carbonate
As can be seen from the equations, only the carbonates produce CO2 giving rise to a pressure increase. However, the water that forms and the water that is present in the HCl solution will also
r 2010 American Chemical Society and Division of Chemical Education, Inc.
_
pubs.acs.org/jchemeduc
nCO2 , pure ¼
_
Vol. 87 No. 12 December 2010
_
msample Mpure
Journal of Chemical Education
ð4Þ
1373
In the Laboratory
Figure 3. Calibration experiment results using CaCO3 of known purity (99.9%). The figure shows the scattering of the test results as a function of sample size.
where msample is the mass of the sample to be analyzed and Mpure is the molar mass of the carbonate in question (e.g., MMgCO3 = 84.32 g/mol). Note that in some cases the dissolution of the sample generates more than 1 mol of CO2 per mole of dissolved carbonate species. In such cases, this has to be accounted for in the equation above with a stoichiometric constant. For example, in the case of pure dolomite, CaMg(CO3)2, 2 mol of CO2 are released for every 1 mol of dolomite dissolved and the right side of the equation above should be multiplied by a factor of 2. The mass and carbonate content of the sample to be analyzed determines the setup requirements. A larger sample will generate more CO2 and require more HCl for a complete reaction. Therefore, it is important to establish an upper boundary for the sample size; it can be calculated based on the stoichiometric reaction ratio between MgCO3 and HCl. For every mole of MgCO3, 2 mol of HCl are required in accordance with eq 1. Therefore, 2 mL of 5 M HCl could dissolve 0.42 g of MgCO3 or 0.50 g of CaCO3. It is also important to know how much CO2 could possibly be generated for reasons of pressure increase and laboratory safety. Dissolving 0.42 g of MgCO3 would result in a significant pressure increase for a small (
