r2010 American Chemical Society and Division of Chemical Education, Inc. _ pubs.acs.org/jchemeduc _ Vol. 87 No. 12 December 2010 _ Journal of Chemical Education 1369 10.1021/ed100353b Published on Web 10/06/2010 In the Laboratory Exploring the Ideal Gas Law through a Quantitative Gasometric Analysis of Nitrogen Produced by the Reaction of Sodium Nitrite with Sulfamic Acid Anne Yu Department of Chemistry, Pomona College, Claremont, California 91711, United States [email protected]The gasometric analysis of nitrogen produced in a reaction between sodium nitrite, NaNO 2 , and sulfamic acid, H(NH 2 )- SO 3 , provides an alternative to more common general chemistry experiments used to study the ideal gas law, such as the experi- ment in which magnesium is reacted with hydrochloric acid and others in which Boyle's law and Charles' law are investigated (1). This experiment, in which the percent sodium nitrite of a sample is determined with the use of equipment commonly found in a general chemistry laboratory, was first published in 1946 (2), and data culled over the last several years show that modifica- tions to the original procedure have improved results in this classic laboratory exercise while also reducing student anxiety related to difficulties in the original procedure. In addition to the ideal gas law, this experiment provides students with an opportunity to analyze data using principles and concepts that constitute core topics often presented in the first semester of general chemistry, such as stoichiometry, redox equations, Dalton's law of partial pressures, the vapor pressure of water, and barometric pressure. The fundamental reaction in this analysis is NO 2 - ðaqÞþ H þ ðaqÞþ NH 2 SO 3 - ðaqÞ f N 2 ðg Þ þ HSO 4 - ðaqÞþ H 2 OðlÞ ð1Þ A 0.25 M sulfamic acid solution is added in excess to samples of an unknown mixture containing sodium nitrite and sodium chloride. The 55-85% sodium nitrite mixtures need to be dried at 110 °C for 1 h prior to the lab period and stored in desiccators as the samples are hygroscopic. For each run, 20 mL of the sulfamic acid solution is required to completely react approxi- mately 0.15 g of unknown mixture. Once the apparatus is assembled and tested for the leaks, the entire reaction proceeds in 20-30 min so that three to five runs may be executed within a 3- to 4-h lab period. Procedure The reaction takes place in a 50 mL Erlenmeyer flask sealed with a no. 2 one-hole rubber stopper with a 24 in. length of 6 mm glass tubing. A 15 in. length of 3/16 in. rubber tubing is attached to the 6 mm glass tubing on one end and the other end of the rubber tubing is attached to a 3 in. length of 5 mm glass inserted into a no. 00 one-hole rubber stopper placed at the top opening of a buret. An additional 20-30 in. of 1/4 in. rubber tubing, the type commonly used for Bunsen burners, is used to connect the tip of the buret to a leveling bulb. A funnel can also be used as a leveling bulb. An 800 mL beaker filled with water is required to provide a constant-temperature bath for the Erlenmeyer flask being used as a reaction vessel. The apparatus is shown in Figure 1, and additional details are provided in the supporting information. Before the reaction initiates and after it is completed, one must ensure that the water in the buret is at the same level as the water in the leveling bulb. This allows the student to employ the barometric pressure as the total pressure of the system. Because the nitrogen is collected over water, its pressure can be calculated after determining the vapor pressure of water at the experimental temperature as the total pressure is the sum of the two. The apparatus must be assembled and checked to ensure that leakage is minimized. The nitrogen gas produced displaces the water in the buret such that the volume can be easily determined using the final and initial volumes. According to eq 1, approximately 0.20-0.25 g of pure sodium nitrite is needed to react completely with 20 mL of 0.25 M sulfamic acid. Approximately 0.14 g of unknown mixture is used for each run; thus, sulfamic acid is present in excess. An easily acquired and inexpensive replacement for the glass vials used in the original experiment (2), which must be cleaned with water, acetone, and briefly oven-dried in between runs, are gelatin capsules. Between the time the capsule is introduced to the acid solution in the reaction vessel and the beginning of the reaction, the student can adjust the leveling bulb to ensure that the level of the volume of the water in the bulb and the buret have equalized. This adjustment often takes several minutes as does Figure 1. The assembled apparatus.
Transcript
r2010 American Chemical Society and Division of Chemical Education, Inc.
_pubs.acs.org/jchemeduc
_Vol. 87 No. 12 December 2010
_Journal of Chemical Education 1369
10.1021/ed100353b Published on Web 10/06/2010
In the Laboratory
Exploring the Ideal Gas Law through aQuantitativeGasometric Analysis of Nitrogen Produced by theReaction of SodiumNitrite with Sulfamic AcidAnne YuDepartment of Chemistry, Pomona College, Claremont, California 91711, United [email protected]
The gasometric analysis of nitrogen produced in a reactionbetween sodium nitrite, NaNO2, and sulfamic acid, H(NH2)-SO3, provides an alternative to more common general chemistryexperiments used to study the ideal gas law, such as the experi-ment in which magnesium is reacted with hydrochloric acid andothers in which Boyle's law and Charles' law are investigated (1).This experiment, in which the percent sodium nitrite of a sampleis determined with the use of equipment commonly found in ageneral chemistry laboratory, was first published in 1946 (2),and data culled over the last several years show that modifica-tions to the original procedure have improved results in thisclassic laboratory exercise while also reducing student anxietyrelated to difficulties in the original procedure. In addition tothe ideal gas law, this experiment provides students with anopportunity to analyze data using principles and concepts thatconstitute core topics often presented in the first semesterof general chemistry, such as stoichiometry, redox equations,Dalton's law of partial pressures, the vapor pressure of water,and barometric pressure.
The fundamental reaction in this analysis is
NO2- !aq"#H#!aq"#NH2SO3
- !aq" f N2!g"#HSO4
- !aq"#H2O!l" !1"
A 0.25 M sulfamic acid solution is added in excess to samples ofan unknown mixture containing sodium nitrite and sodiumchloride. The 55-85% sodium nitrite mixtures need to be driedat 110 !C for 1 h prior to the lab period and stored in desiccatorsas the samples are hygroscopic. For each run, 20 mL of thesulfamic acid solution is required to completely react approxi-mately 0.15 g of unknown mixture. Once the apparatus isassembled and tested for the leaks, the entire reaction proceedsin 20-30min so that three to five runs may be executed within a3- to 4-h lab period.
Procedure
The reaction takes place in a 50 mL Erlenmeyer flask sealedwith a no. 2 one-hole rubber stopper with a 24 in. length of 6mmglass tubing. A 15 in. length of 3/16 in. rubber tubing is attachedto the 6 mm glass tubing on one end and the other end of therubber tubing is attached to a 3 in. length of 5 mm glass insertedinto a no. 00 one-hole rubber stopper placed at the top openingof a buret. An additional 20-30 in. of 1/4 in. rubber tubing, thetype commonly used for Bunsen burners, is used to connect thetip of the buret to a leveling bulb. A funnel can also be used as aleveling bulb. An 800 mL beaker filled with water is required
to provide a constant-temperature bath for the Erlenmeyerflask being used as a reaction vessel. The apparatus is shown inFigure 1, and additional details are provided in the supportinginformation.
Before the reaction initiates and after it is completed, onemust ensure that the water in the buret is at the same level as thewater in the leveling bulb. This allows the student to employ thebarometric pressure as the total pressure of the system. Becausethe nitrogen is collected over water, its pressure can be calculatedafter determining the vapor pressure of water at the experimentaltemperature as the total pressure is the sum of the two. Theapparatus must be assembled and checked to ensure that leakageis minimized. The nitrogen gas produced displaces the water inthe buret such that the volume can be easily determined using thefinal and initial volumes.
According to eq 1, approximately 0.20-0.25 g of puresodium nitrite is needed to react completely with 20 mL of0.25M sulfamic acid. Approximately 0.14 g of unknownmixtureis used for each run; thus, sulfamic acid is present in excess. Aneasily acquired and inexpensive replacement for the glass vialsused in the original experiment (2), which must be cleaned withwater, acetone, and briefly oven-dried in between runs, aregelatin capsules. Between the time the capsule is introduced tothe acid solution in the reaction vessel and the beginning of thereaction, the student can adjust the leveling bulb to ensure thatthe level of the volume of the water in the bulb and the buret haveequalized. This adjustment often takes several minutes as does
Figure 1. The assembled apparatus.
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In the Laboratory
the time required for the acid to dissolve through the capsule.Consequently, the delay in the reaction that is inherent in usinggelatin capsules allows for the experiment setup to be executedmore precisely.
Hazards and Disposal
Students should wear eye protection and lab coats.Sulfamic acid should be handled carefully to avoid spillingand contact with the skin. Spills can be neutralized withsodium bicarbonate, which should be readily available in thelab. Sodium nitrite is not combustible, but is a strong oxidizerand, similar to sulfamic acid, is harmful if swallowed and canbe irritating to the skin, eyes, and respiratory tract. Contactwith the skin and eyes should be followed with flushing withwater. If fresh air does not improve breathing after inhalation,then seek medical attention. The instructor should cautionstudents that when cleaning up after each run to dispose of thegelatin capsule remainders into the trash, rather than downthe sink.
Data Analysis
The data collected are the barometric pressure, temperature,volume change, and initial mass of the sodium nitrite unknown.The pressure of the gas is experimentally made equivalent to thebarometric pressure and can be calculated employing Dalton'slaw of partial pressures after determining the vapor pressure ofwater at the experiment temperature. The student is then able toemploy the ideal gas law
PV ! nRT "2#
to determine the amount of nitrogen gas evolved. From this thepercent composition of sodium nitrite in the initial sample canbe determined from the stoichiometric relationship given ineq 1.
Results
Three years of student lab results are shown in Table 1.Each student reported the average of at least three trials. Thetable lists the averages and standard deviations of all studentresults for each of the unknown mixtures. The original articleclaims that an accuracy of 0.1 to 0.2% may be expected (2), butthe student data show otherwise. There is an expected errorranging from 1-5% as a result of several inherent experimentalerrors. For example, a 1 !C rise in room and bath temperaturewould result in a 0.5% error. The solubility of nitrogen gas maycause another 0.15% error for a sample that generated a volume
of 40 mL. The reaction of empty gelatin capsules with the acidsolution showed no contribution to the volume of gas pro-duced.
The error increases as the percent composition of sodiumnitrite increases in the unknown sample. This is likely due to theincrease in reaction time. As more time is required to allowcomplete evolution of gas, there is also an increase likelihood ofleakage. This would decrease the volume change and thus theamount of sodium nitrite calculated. Additional results similar tothose reported in Table 1 from previous years of using glass vialscan be provided upon request. The results indicate that there isan overall improvement in the average percent compositionusing gelatin capsules with an overall smaller range of standarddeviations. For optimal precision and accuracy, unknown sam-ples that contain approximately 55-70% sodium nitrite workbest. The data in Table 1 are helpful in making peer comparisonsshould an instructor distribute a range of samples among thestudents.
In the original experiment, an open glass vial was used, andupon contact with the sulfamic acid, the reaction proceededimmediately. It was common for the reaction to start inadver-tently before students had time to equalize the water levels dueto accidental tipping over of the open sample vial in the beakercontaining sulfamic acid. This would have led to misleadingvolumes and greater deviations in the averages if runs hadinitiated too early and on an inconsistent basis before waterlevels had been properly adjusted. The inadvertent tipping overof the open glass vials was an enormous source of frustration,and replacement of the glass vials with gelatin capsule hasremoved the anxiety associated with the experiment, increasedefficiency during the lab period, and reduced the waste ofchemicals.
Summary
The gasometric analysis of nitrogen produced in thereaction of sodium nitrite with sulfamic acid is rich inchemistry and can be executed largely with items commonlyfound in a general chemistry laboratory. The experimentallows students to employ a variety of concepts presented inthe lecture portion of the curriculum as they analyze data. Thequantitative nature of the experiment is increasingly moreprecise and accurate with replacement of the glass vials withgel capsules as it provides more time to allow equalization ofpressure. Though there are inherent and unavoidable errors,the 50-70% sodium nitrite samples have resulted in betteraccuracy, requiring less time to completely react than thehigher sodium nitrite content. Once the apparatus is set up,students are able to perform multiple runs in a 3- or 4-h lab
Table 1. Comparison of Results Using the Glass Vials and Gelatin Capsules
Results with Glass Vial Results with Gelatin Capsule
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In the Laboratory
period, which should yield ample data for a quantitativeanalysis of the results.
Acknowledgment
The author thanks the members of the Pomona CollegeChemistry Department who were initially involved in imple-menting this experiment, especially Chuck Taylor for sugges-tions and discussion, Elena Branford (class of 2010), who testedthe modification, and the general chemistry students whoseresults are included here.
Literature Cited1. Wilbraham, A.; Staley, D.; Simpson, C.; Matta, M. Chemistry
Laboratory Manual, 2nd ed.; Addison-Wesley: New York, 1990;pp 135-142.
2. Brasted, R. J. Chem. Educ. 1946, 23, 320–321.
Supporting Information Available
Instructions for the students; notes for the instructor. Thismaterial is available via the Internet at http://pubs.acs.org.
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Gas Analysis of Nitrite 1
LAB DOCUMENTATION for the student experiment
GAS ANALYSIS OF NITRITE
The objective of this experiment is to determine the percent of NaN02 in a
sample by measuring the volume of N2 gas liberated by the action of an
excess of sulfamic acid, H(NH2)SO3, and to gain an understanding of the
applications of the gas laws.
PRINCIPLES
The fundamental reaction in this analysis is
NO2-
(aq) + H+(aq)+ NH2SO3
-(aq)! N2 (g) + HSO4
-(aq)
+ H2O (l) (1)
Although the reaction is specific for NO2
-, the methods employed are very general and
could be applied to the analysis of any relatively rapid reaction in which a gas is evolved
as the result of mixing a liquid reagent with another solid or liquid. The method cannot
be applied (a) if the gas produced is very soluble in the liquid solution, (b) if the reaction
does not go to completion, or (c) if two or more gases are formed simultaneously and in
variable relative amounts.
This experiment also illustrates several different applications of the gas laws as well as
some aspects of the mechanical behavior of confined gases which are not always apparent
to the novice. The fundamental principles are discussed in your text, with the "ideal gas
law"
PV = nRT (2)
being basic to the entire discussion. In brief outline, chemical Eq. (l) shows that for every
mole of N2 (g) produced, one mole of NO2- must have been present in the original
unknown sample. Eq. (2) shows that the number of moles of N2(g) produced can be
obtained by measuring the volume of N2(g) at the measured conditions of temperature and
pressure. In this experiment, the total pressure (equal to barometric pressure, PB) is
measured. Because the N2 is evolved from an aqueous solution it is saturated with water
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Gas Analysis of Nitrite 2
vapor, and the total pressure is the sum of the pressure of the N2 gas that is produced and
the vapor pressure of water with which it is saturated. As a consequence, you must apply
Dalton's Law of partial pressures in order to obtain the true pressure of the N2 alone in the
Since the measuring system is closed, the N2 is saturated with water vapor and PH2O will
be equal to the vapor pressure of the liquid water at the temperature of the experiment.
The barometric pressure, PB, can be read directly from the laboratory barometer.
Since small changes in temperature during the course of a given run can have a serious
effect, every effort must be made to keep the temperature constant. For example, if there
were 60 ml of trapped air in the closed system, a 1°C rise in room and bath temperature of
27°C would cause a change of (1/300) (60) = 0.2 ml in the water level of the buret, even
though no N2 had been generated. If 40 ml of N2 gas were actually generated, this error
would correspond to a percentage error of 0.5%. The other major error, but one over
which you have no control, is the solubility of the nitrogen (and therefore the loss of the
N2) in the reaction mixture. This solubility, over and above that already dissolved from
the air, which is 79% N2, is approximately 0.06 ml, an amount corresponding to about
0.15% error for a N2 volume of 40 ml. These and other small inherent errors combine to
make gas analysis, as illustrated by the simple apparatus of this experiment, appreciably
less precise than many of the other methods studied in this course. It is reasonable to
expect an overall relative error of the order of one percent.
EXPERIMENTAL PROCEDURE
1. Special Equipment Required
You will each need the following equipment:
(a) 1 leveling bulb
(b) 1 No. 00 one-hole rubber stopper with a 3-in. length of 5-mm glass tubing
(c) 1 No. 2 one-hole rubber stopper with a 24-in. length of 6-mm glass tubing
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Gas Analysis of Nitrite 3
(d) 1 15-in. length of 3/16-in. rubber tubing
2. Apparatus Assembly
(a) Assemble the gas analysis apparatus according to the following figure. The two
rubber stoppers should be slightly moistened so as to insure a gas-tight seal. Attach a
clamp to a ring stand for the purpose of holding the leveling bulb. Fill your 800-ml
beaker about three-fourths full with tap water so that it can be set aside and come
to room temperature while you are preparing for the experiment.
(b) Remove the stopper from the top of the buret and slowly pour de-ionized
water into the buret so that when the water level in the buret is just below the
zero mark, the other water level stands at the lower end of the leveling bulb.
Make certain that no air bubbles are trapped in the connecting tubing by
running the water back and forth between the bulb and the buret. Replace the
Our experimental set-up is similar to a gas manometer. Instead of measuring a height difference in the two arms, however, we adjust the height of the leveling bulb so that h = 0, and Pint = Patm. For this reason, all volume measurements must be made when the water levels are equal.
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Gas Analysis of Nitrite 4
stopper at the top of the buret while the bulb is clamped level with the top of
the buret.
(c) Test for leaks by lowering the leveling bulb and observing whether, after the
initial drop in the water level in the buret, the level stays in a fixed position or
drops over a period of time due to a leak. If necessary, find and eliminate the
leak.
(d) Remove the 50-ml Erlenmeyer flask while the leveling bulb is level with the
top of the buret.
3. Weighing the Sample. Obtain an unknown sample from the desiccator at the side
shelf and transfer it immediately to your own desiccator. Record the sample number.
Calculate in advance the approximate weight of sample needed, assuming that the sample
contains about 70% NaNO2, you would like to collect about 35 ml of nitrogen gas, the
pressure is 710 torr, and the temperature is 22˚C. Tare a gelatin capsule on the balance.
To prevent picking up any residue from the balance, place the capsule on a clean watch
glass. Then quickly, but carefully transfer the pre-calculated amount of sample of the
unknown from your desiccator to creased weighing paper tared on the balance. Use the
crease in your weighing paper to quickly make a transfer of the unknown into the capsule.
Reweigh the enclosed unknown sample in the capsule to accurately determine the mass
that will be reacted. The samples are quite hygroscopic so keep your supply of unknown
in the desiccator when not needed for weighing. It will not affect your results if your
weighed sample picks up moisture after it has been weighed.
4. Preparing the Reaction Flask
(a) Use your 10-ml pipet to add approximately 20 ml of 0.25 M sulfamic acid
solution to the flask.
(b) Add your capsule with unknown into the flask. You will have about five
minutes before the capsule begins to dissolve to complete steps 4(c) – 5(b).
(c) Moisten the rubber stopper and carefully re-connect the Erlenmeyer flask. Bring
the 800-ml beaker of water up around the flask and support it with the tripod.
Place the leveling bulb near the buret and adjust the bulb up or down so
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Gas Analysis of Nitrite 5
that the water levels in the bulb and buret are at the same height. Wait a
few minutes to make certain temperature equilibrium is reached so that the
levels do not change.
5. Carrying out the Reaction
(a) Record the temperature of the water in the beaker and the air temperature in the
room, estimating it to the nearest 0.1°C. Record the water level in the buret,
estimating it to the nearest 0.01 ml.
(b) Lower the leveling bulb. Unclamp the glass tube holding the Erlenmeyer flask in
the water, remove the flask from the beaker, and then shake it gently. Replace the
flask in the beaker and again clamp the glass tube.
(c) Adjust the bulb so as to keep the two water levels at roughly the same height;
periodic readjustment will be needed for perhaps 15 minutes or more until the
capsule has dissolved enough to allow for all of its contents to react. It is
important to shake the reaction flask vigorously several times during this time to
make sure that the reaction has gone to completion. During this waiting period
you could also weigh out another sample of unknown in another vial, following
the precaution noted in 3.
(d) Continue to monitor the level of the water in the bulb and the buret after the
reaction has gone to completion. Adjust the level of the bulb so that the water
remains at the same height as it is in the buret. This may take a few minutes.
When the water levels remain unchanged for a few minutes, record the new water
level in the buret to the nearest 0.01 ml. Record also the temperature of the air
surrounding the buret to the nearest 0.1°C. When no further change in levels is
apparent, make sure that the temperature of the water in the beaker is identical to
that observed in 5(a). If the water temperature has changed in the beaker, stir the
water and cool (with ice) or heat (with burner) as needed. Do this with care, for it
is very easy to overdo it.
6. Repeating the Determination
(a) Remove beaker, disconnect the Erlenmeyer flask, and remove solution and vial.
Rinse the flask and dry the neck to prevent drops from contaminating your next
sample during its insertion into the flask.
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Gas Analysis of Nitrite 6
(b) Repeat the experiment as many times as desired (three minimum), following the
procedures 3-5. You may wish to adjust the sample size in succeeding runs to
give you a conveniently larger or smaller volume of gas.
(c) Determine the barometric pressure (PB) from your reading of the laboratory
barometer. If you are collecting data over more than one lab period, make sure to
read the pressure each day.
CALCULATIONS
1. Calculate the weight percent of NaNO2 in the original sample. Make sure
to take into account the vapor pressures of water, which can be looked up in a
reference manual, such as the CRC.
2. Calculate the mean value of your results, the standard deviation, and the
relative 95% confidence interval of the mean.
3. Hand in the report sheet summarizing your results.
4. Be sure to show your sample calculations and to answer the question on
the report sheet.
PRE-LAB 1) Calculate how many grams of unknown nitrite sample to use in each trial. Assume your sample is 70.0 % NaNO2 and that the reaction produces a volume of 35.00 mL at 730 torr and 22.0 °C? Remember to account for the vapor pressure at this temperature. 2) Suppose your unknown sample were 60.0% nitrite. Would the volume produced be higher or lower compared to the volume produced under the same conditions given in question #1? Would you need more or less grams of unknown to obtain a 35 mL volume change?
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Gas Analysis of Nitrite 7
Name _________________________________ Lab Section __________________ Date Report Submitted ________________ Sample No. __________________
GAS ANALYSIS OF NITRITE
1. Percentage NaNO2 in Sample First Lab Day Second Lab Day Observed barometric pressure, PB, mm __________ __________
5 _______ _______ _______ _______ _______ _______ _______ Average _____________ Standard deviation _____________ Relative 95% confidence interval of the mean, % _____________
Indicate which runs are used to calculate the average by circling their numbers.
2. Show sample calculations on the reverse side.
3. Solve the following problem. Show the calculations on the reverse side.
For your first sample reported above how many moles of water vapor were contained in your reported gas volume?
Moles of water vapor ____________
1
LAB DOCUMENTATION - the instructor’s notes
GAS ANALYSIS OF NITRITE-Instructor’s Notes
Before the demonstration have students fill their 800 ml beakers with tap water and allow to
equilibrate to room temperature. This is the constant temperature bath for the reaction.
Answers to pre- lab questions: 1. n (N2) = PV/RT = [(730-19.8) /760 ]atm x 0.0035L = 0.9344 x .035 =.000135 0.08206 Latm/mol K x (273+22) K 24.2077 0.00135 mol of N2 = 0.00135 mol of sodium nitrite ! 0.00135 moles of sodium nitrite x 69 g/mol ! 0.09315 g of sodium nitrite 0.09315 g of sodium nitrite = .70 x (Y g of unknown) ! Y = 0.1331 g of unknown 2. The volume would be lower, requiring more grams of unknown.
Notes on experiment:
1. Discuss briefly the chemistry of this experiment and how the experiment illustrates the ideal gas law
and Dalton’s Law of Partial Pressures.
NO2-(aq) + H+
(aq) + NH2SO3-
(aq) ! N2(g) + HSO4-
(aq) + H2O(l)
PB = PN2 + PH2O
2. Sample vials of unknown are obtained from a stock desiccator and immediately put into a transport
desiccator. NaNO2 samples are very hygroscopic. The sample should be weighed quickly and the
rest of sample in the sample vial returned to the transport desiccator immediately. Weigh only 1
sample at a time. At the end of the day, sample vials must be returned to the stock desiccators. A
sample vial cannot be stored in a transport desiccator from one period to the next period.
3. For estimating the size of sample to use, assume 70% NaNO2, 730 torr, 22.0 degrees C and a volume
of 35 ml. This was one of their pre-lab questions. After lst run adjust sample size to get 30-40 ml
volume change, if necessary.
4. Demonstrate and discuss the operation of the set-up.
PH2O determined from reference source, such as CRC; note that students should interpolate between tabulated values to get value at measured T.
2
a. Mention need to moisten stoppers and "twist," not push all connections. Be careful not to
break buret etc.
b. Fill buret with de-ionized water. Make sure there are no bubbles or leaks.
c. Show how easily to get and keep air bubbles out of leveling bulb tubes. Fill leveling bulb
with de-ionized water while pinching off the end of the rubber tubing. Then, while holding it
over the sink, suddenly let out a big slug of water. Finally, while it is still pinched off,
connect the tubing to the tip of the buret. The buret should already be full of de-ionized
water, but level below zero mark. Massage the rubber tubing to get rid of air bubbles,
especially where the tubing covers the tip of the buret.
d. Students should use LATEX tubing not pressure tubing with the leveling bulb.
e. Show how to test for leaks by attaching empty Erlenmeyer to close the system and then
lowering leveling bulb and reading buret. Don't forget to open stopcock! The stopper can
pop off if the gas buildups while the stopcock is shut. Wait 5 minutes and read again—level
should not change. This part is important because all of your results will be affected if there
is a leak. If necessary, find and eliminate leak. Be sure to remoisten Erlenmeyer stopper
each time you re-connect.
f. Demonstrate the various ways to adapt system to individual tripods: tall stands over buret
stand base; short ones to the side of the base.
g. Show how to fix leveling bulb ring stand for greatest flexibility and easiest matching of
water levels. You need to use footstands to substitute for the hoods in the lab.
h. 20 mL acid delivered via Mohr pipet (Discuss use of Mohr pipet including how to read
markings) in previous years, but the modified lab no longer requires the use of a Mohr pipet.
Pour approximate amount of sulfamic acid into beaker and pipet from there to avoid
contamination of stock solution.
I think the best way to go about a transfer into a capsule is to weigh an empty capsule on a
watchglass or weighing paper so it does not pick up any residue from the balance. Add a sample
to the weighing paper and transfer carefully to the capsule. Then reweigh the capsule, making
sure that no residue of the nitrite is on the outside of the capsule. It can wiped off with a
Kimwipe.
i. Before connecting the Erlenmeyer with the sample, raise leveling bulb until levels equalize
with the buret reading near 0 mL and only 2-3 cm of water showing in the leveling bulb.
Connect the Erlenmeyer carefully but firmly. Adjust buret and bulb water levels until they
are the same. Read buret volume. Measure the temperature of the bath and the temperature
of air near the top of the buret (for PV = nRT). Make sure thermometer is dry for the air
measurement. Lower leveling bulb before starting reaction.
3
j. Show how to hold flask while vigorously shaking it for complete evolution of gas; care with
connections, etc.; repeat shaking several times during course of reaction to ensure
completeness. Reaction takes 20-25 minutes. (During this time, students can measure out
another sample and/or prepare additional tables in their notebooks for the next run, etc.)
k. Re-measure the air and bath temperatures at the end of the reaction. The temperature of the
bath should be within ± 0.1 degree C of the initial temperature. Ice will be available in the
lab to adjust the beaker temperature when necessary.
l. Read buret volume at end of reaction with buret and bulb levels the same.
5. Show the students how to use the electronic barometer and where it will be in lab.
6. For estimating precision and for making the decision about whether additional runs should be made,
compare the ratios of observed volume of N2 to weight of sample, assuming that the pressure and
temperature for the runs being compared are the same. For best results, three successive runs within
l% relative range should be accomplished. (Should have time to do 2 or 3 runs the first day, and as
many as 4 or 5 the second day if necessary.)
Preparation: Mixtures of NaCl and NaNO2 can be mixed in different percentages by mass. 50-75% works better than higher percents. Mixtures are milled for 2hours to ensure homogeneity. Oven dry unknowns at 110 oC for 1 hour. Keep stored in a dessicator. Each student will need about 0.20 grams of unknown for each run required. 0.25 M sulfamic acid (H2NSO3H) can be made by dissolve 24.5 grams anhydrous sulfamic acid in 1 liter of water. Keep in colored glass bottles. Each student will need 20 mL of acid for each run required. CAS Registry Numbers: NaCl 7647-14-5 NaNO2 7632-00-0 H2NSO3H 5329-14-6 Other Material(s): Gel Capsules – size 00 ordered from SPI Supplies (#02302-SS)
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Gasometric Determination of CO2 Released fromCarbonate MaterialsJohan Fagerlund* and Ron ZevenhovenThermal and Flow Engineering Laboratory, Åbo Akademi University, Turku, Finland*[email protected]
Stig-G!oran Huld"en and Berndt S!odergårdCombustion and Materials Chemistry, Åbo Akademi Process Chemistry Centre, Turku, Finland
As part of the ongoing mineral carbonation research, aninexpensive and simple method for determining the carbona-tion content of various carbonated materials (mainly carbo-nated brucite, the mineral form of magnesium hydroxide) wasrequired. Mineral carbonation is part of a larger carbon dioxidecapture and storage (CCS) portfolio and has received com-paratively little attention despite its inherent potential(natural process, abundant resources, and permanent CO2storage). For more information about this area, see, forexample, the latest literature reviews (1, 2).
The method described here is of value to mineral carbona-tion research owing to its simplicity, low cost, and accuracy, but itcould also serve another purpose; that is, to expose undergraduatestudents to a hands-on implementation of the ideal gas law andsome basic inorganic chemistry. In addition, we consider aspectsthat could easily be overlooked in a theoretical approach, such asthe influence of the heat generated during the reaction between asolid carbonate and hydrochloric acid. In a classroom, theseaspects could be considered beforehand, for example, in groups,but it could also be useful to allow the students to find theseparameters by simple trial and error. For instance, one studentmight take the sample container into his or her hand withoutrealizing how much it affects the temperature inside the containerbefore he or she compares the obtained result with that ofothers.
The method presented here is based on the reactionbetween a carbonated material and an aqueous solution ofHCl, which generates CO2. This method is not new andliterature dating back to the late 1940s has described thismethod (3). The reaction system is sealed and the pressureincrease is used as a measure of the carbonation degree of thesample, as has also been noted by others more recently (4-6).Choi andWong used a similar method to determine the reactionkinetics and carbonate content of eggshells, but did not assess theaccuracy of the method (4). Roser andMcCluskey demonstratedthe use of pressure measurement and data logging to determinethe stoichiometry of reactions between various carbonates andHCl (5). Horvath et al. was able to accurately determine thecarbonate content of soils measuring the pressure increase insidea sealed vessel (6). The method described here takes a slightlydifferent approach, combining several aspects not described inany single literature source cited above.
Pile et al. from the Los Alamos National Laboratory(LANL) developed a method for determining the carbonatecontent of mineral carbonates (7). They measure the carbonatecontent indirectly through volume change and hydrostatic
pressure; however, we measure the pressure directly. Despitethe simple design and reported high accuracy of the LANLsystem, we were unable to accurately repeat the tests onmaterials that may be low in carbonate and high in magnesiumhydroxide and decided to use (similarly to refs 4-6)) themethod described here.
Design and Construction of Apparatus
The equipment used to determine the CO2 release of amaterial is simple and available in most teaching laboratories, themost important component being a pressure gauge. Any pressuremeasurement device capable of determining the pressure to aprecision of about 100 Pa could be used. In principle, a manometercould suffice, although we used a PASPORT Absolute Pressure/Temperature Sensor, PS-2146. Together with PASCO's softwarecalled DataStudio, the pressure (and temperature) can be loggedand displayed simultaneously as the experiment proceeds. Thisequipment is especially developed for educational purposes. Inaddition to a pressure gauge, a container for the sample, a smallercontainer for the acid, an airtight plug, and some flexible tubes toattach to the pressure gauge are needed. The container for thesample needs to be durable and transparent, such as a sturdy glass jaras 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, itis preferable for more consistent results. The system should notbe changed after it has been assembled as any change requires thesystem to be recalibrated (see the calibration section). Thetemperature sensor should be located inside the glass jar. It isalso possible to assume constant temperature, but this gives lessaccurate 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 theglass jar. The carbonate sample and theHCl container are carefullyinserted into the glass jar together with a magnetic stir bar. Theglass jar is sealed with a rubber cap containing the temperaturesensor and a tube that is connected to a pressure gauge. The datalogger (if available) is started and the system is stabilized (humiditybuilds up) for a short period of time. The magnetic stirring deviceis started, which spills the HCl solution, and the reaction is left togo to completion, that is, until no more pressure increase isobserved.
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The experiment takes less than 10 min, but it is importantthat the system is allowed sufficient time to reach thermody-namic and chemical equilibrium before and after the stirring,ensuring that all carbonates have dissolved. The steady state ofthe system can be determined using the pressure gauge. If thepressure is stable, so is the system. It is helpful for the students tosee the progress of the reaction on a computer screen (examplesin Figure 4).
Hazards
Hydrochloric acid is corrosive and causes burns to all bodytissue.
Calculations
Using the ideal gas law with a constant volume system, one ofremaining three variables can be determined knowing the othertwo variables. In this experiment, the temperature and the pressureare measured so the only unknown variable is the amount of thesubstance. Assuming that the created atmosphere consistsmostly ofCO2 and a small quantity of water, the mass of CO2 released fromthe sample can be calculated. It is also worth mentioning that thedissolution of CO2 into the solution is negligible owing to the highacidity of the solution and the magnetic stirring. The reactiontaking place inside the reaction vessel is
MgCO3!s"# 2HCl!aq" f Mg2#!aq"# 2Cl- !aq"#H2O!aq"#CO2!g" !1"
Similarly, knowing that most of the samples also contain magne-sium hydroxide or magnesium oxide, the reaction equations forthese with HCl are
MgO!s"# 2HCl!aq" f Mg2#!aq"# 2Cl- !aq"#H2O!aq"!2"
Mg!OH"2!s"# 2HCl!aq" f Mg2#!aq"# 2Cl- !aq"# 2H2O!aq" !3"
As can be seen from the equations, only the carbonates produceCO2 giving rise to a pressure increase. However, the water thatforms and the water that is present in the HCl solution will also
increase the system pressure to a small degree owing to theinevitable vaporization of some of the water (HCl vaporizationcontributes only a very small fraction of the water partial pressureand is therefore neglected, see the supporting information andref 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 canbe shown (Figure 2) that most of the water evaporating from theHCl solution takes place before the reaction (typically around 500 s)starts. Allowing the system to stabilize before the reaction elim-inates 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 ofMgO andMg(OH)2 is more exothermic (9.3$and 6.7$, respectively) than that of MgCO3; thus, more heat isreleased when either MgO or Mg(OH)2 is in the sample. Thiseffect can clearly be seen from the data in Figure 2. In the case of thesample containing onlyMgCO3, the relative humidity increase dueto the dissolution reaction is barely visible (around 1000 s), whereasin the case of a mixture of MgCO3 and Mg(OH)2, the increase inboth temperature and relative humidity is obvious (around 1600 s).In other words, water vapor plays a less significant role in sampleswith a high carbonation degree, whereas samples with a lowercarbonation degree can easily be overestimated. The relativehumidity 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 theamount of CO2 released from the sample to the calculated releasefrom a 100% pure carbonate of the same species and mass. Theamount of CO2 in a 100% pure carbonate sample is calculatedknowing the molar mass of the carbonate
nCO2, pure %msample
Mpure!4"
Figure 1. Experimental setup.
Figure 2. Relative humidity as a function of time inside the glass jar fortwo different experiments. Note the different time scales of the twoimages.
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wheremsample is the mass of the sample to be analyzed andMpureis the molar mass of the carbonate in question (e.g., MMgCO3
=84.32 g/mol). Note that in some cases the dissolution of thesample generates more than 1 mol of CO2 per mole of dissolvedcarbonate species. In such cases, this has to be accountedfor in the equation above with a stoichiometric constant. Forexample, in the case of pure dolomite, CaMg(CO3)2, 2 mol ofCO2 are released for every 1 mol of dolomite dissolved and theright side of the equation above should be multiplied by a factorof 2.
The mass and carbonate content of the sample to beanalyzed determines the setup requirements. A larger samplewill generate more CO2 and require more HCl for a completereaction. Therefore, it is important to establish an upperboundary for the sample size; it can be calculated based on thestoichiometric reaction ratio between MgCO3 and HCl. Forevery mole of MgCO3, 2 mol of HCl are required in accordancewith eq 1. Therefore, 2 mL of 5 M HCl could dissolve 0.42 g ofMgCO3 or 0.50 g of CaCO3. It is also important to know howmuch CO2 could possibly be generated for reasons of pressureincrease and laboratory safety. Dissolving 0.42 g of MgCO3would result in a significant pressure increase for a small (<200 mL)system volume.Assuming atmospheric pressure in the beginning andconstant temperature throughout the reaction, the final pressurewould become 167.6 kPa for a 150 mL system and 199.3 kPa for a100 mL system. Taking into account a relatively large temperatureincrease, for example, from 21 to 24 !C, the final pressure would notchange much, increasing only 1.7 kPa for the 150 mL system and2.0 for the 100 mL system.
Knowing the sample and the maximum amount of CO2that a sample could generate allows for calculation of thecarbonation degree. The carbonation degree is defined accordingto the following formula:
! !nCO2, gennCO2, pure
! nfinal - ninitnCO2, pure
"5#
where nCO2,gen is the amount of CO2 released or generatedduring an experiment and is the difference between the moles ofgas (air$ generated CO2) at the end, nfinal, and at the start (onlyair), ninit. Thus, the carbonation degree is determined knowingonly the initial and final conditions inside the sample holder (orglass jar). The most important aspect of this method is thereforeto accurately measure the pressure before any reactions inside theglass jar take place and after, when all the reactions have stopped.In an ideal case, the temperature should be the same before and
after the experiment, but in practice, this would require longstabilization times.
In our experiments, we noticed that the temperature insidethe glass jar increases owing to the exothermic nature of thereactions and it takes a long time before the temperature returnsto the initial value. Therefore, to speed up the process, thetemperature reading can be taken before the system has reachedsteady state with its surroundings if a temperature sensor isplaced inside the glass jar.
Before any reaction has taken place in the reaction vessel, theinitial amount of gas (air) can be calculated using the ideal gaslaw. Similarly, the final amount can be calculated knowing thefinal pressure and temperature
ninit !pinitVsystem
RTinit"6#
nfinal !pfinalVsystem
RTfinal"7#
where pinit and pfinal are the pressure before and after the reaction,respectively; Vsystem is the total volume of the system; R is theuniversal gas constant; and Tinit and Tfinal are the temperaturesbefore and after the reaction, respectively. Vsystem is a term thatcould be measured or approximated knowing the dimensions ofthe parts making up the experimental setup, but it is easier tosimply obtain this value using a sample of known carbonatecontent (e.g., 99.9% pure CaCO3). The calibration of the systemis only required the first time or when the setup has been altered.
Calibration Experiment
System VolumeTo determine the system volume (including the glass jar and
tubes and excluding measurement instruments), a calibrationexperiment is necessary. A calibration experiment is done using acalcium carbonate sample of a known carbonation degree.Measuring the temperatures and pressures before and after thereaction allows the determination of Vsystem using
Vsystem !R!
msample
Mpure
!
pfinalTfinal
! "-
pinitTinit
! " "8#
Equation 8 was derived by combining eqs 4, 6, and 7 and solvingforVsystem (for details see the supporting information). Note thatCO2 is not an ideal gas, but under the pressure and temperaturerange investigated, this is a valid assumption.
System ErrorUsing the average calculated volume, the data were re-
worked to obtain the carbonation degree of the known CaCO3sample. The results using 99.9% pure CaCO3 are displayed inFigure 3. The average of 38 samples is 99.97% and the standard(and relative standard) deviation is 0.9%. Two of the 38measurements were excluded as outliers based on Peirce's criter-ion (9). In an ideal case, all points would lie on the samehorizontal line, which would mark the 99.9% carbonationdegree. However, in practice, there are always errors, seen asthe random scattering of points around the 99.9% line in
Figure 3. Calibration experiment results using CaCO3 of known purity(99.9%). The figure shows the scattering of the test results as a function ofsample size.
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Figure 3. More information and details about the calibrationcalculation procedure can be found in the supporting information.
Experiment Data
For the calculations, we only need two data samplings duringan experiment, one at the start and one at the end. However, it isalso important to observe the temperature and pressure trendsobtained by continuous sampling, especially to assess stabiliza-tion. A number of points for the temperature and pressure duringthree typical carbonate analysis experiments are shown in Figure 4.Notice that the ambient pressure can be obtained at time zero andthat shortly after this the rubber cap is inserted into the glass jar(point A) resulting in a small pressure peak. In addition, thetemperature rises a little owing to compression of the gas inside theglass jar, and because of the aqueous HCl solution, the humidityinside the glass jar will also increase (see the supportinginformation). This can be seen as a slight increase in pressure afterpointA.The pressure stabilizes quickly, but usually the temperaturetakes a little longer to stabilize. At this point, the magnetic stirrercan be turned on (or the flask can simply be shaken by hand)starting the reaction. The reaction is very fast as can be seen fromthe steep increase in pressure at point B (the temperature followswith a small delay). This increase in pressure is the result of CO2formation in the system and can be used to determine thecarbonation degree of the sample.
The graphs of three experiments using different samples areshown in Figure 4. The top graph was obtained using a mixtureof 0.0963 g of MgCO3 and 0.0982 g of Mg(OH)2, the middlegraph was obtained using 0.1555 g ofMg(OH)2, and the bottomgraph is the result of an experiment without a sample, just theHCl solution. The values required for the calculations of thecarbonate content in Figure 4 are taken just before the stirring(point B) and sometime after the peak pressure has been reached(point C). The exact position of point C is not critical as long asthe temperature is recorded at the same time. However, it ispreferable to allow the temperature to stabilize for a while to obtaina representative temperature value for the entire system. The largedrop in both temperature and pressure (point D) is due to theremoval of the rubber cap sealing the glass jar.Note the small increasein pressure for the “no solid sample” experiment; this is caused by theincrease in vapor pressure initialized by the stirring. The increase inpressure for theMg(OH)2 experiment is caused by the small quantityof impurities (3.8%-wt CO3) in the sample.
Verifying the Method
For the purposes of determining the carbonation degree of asample quickly and without expensive equipment, the methodproposed here performs very well. The data in Figure 3 show thatmost of the experiments lie within a margin of (1% from thetarget line. However, this is for calcium carbonate and we aremore interested in magnesium carbonate/hydroxide mixtures.Therefore, the method was also tested on a different set ofsamples. The testing on samples consisted of a mix of SiO2(sand), commercial brucite (Dead Sea Periclase Ltd.), andmagnesium carbonate (Fluka). The materials also containedsome impurities, so the carbonate content of the brucite andmagnesium carbonate was tested in a certified laboratory forreference purposes. The commercial brucite contained 38 gCO3/kg and the magnesium carbonate had 461 g CO3/kg.
The results of the five tests are given in Figure 5 showing thecarbonation degree based on the certified laboratory results onthe y axis versus the measured carbonate content on the x axis. Inthe best case, the dotted line in Figure 5 should pass perfectlythrough all of the points and y would equal x. This is nearly thecase and the error is small (R2 = 0.995) showing that the methodcan be used to accurately find out carbonate contents of solidmaterials. However, it should be noted that not all materials canbe used. Some materials could react to form other gaseouscomponents besides CO2 and H2O, causing the carbonatecontent of the sample to be overestimated.
Sources of Error
Knowing how to solve a problem in theory is a completelydifferent than solving it in practice. Although the experiment
Figure 4. Temperature and pressure profiles of carbonate analysisexperiments: (top) a mixture of 0.0963 g of MgCO3 and 0.0982 g ofMg(OH)2; (middle) 0.1555g ofMg(OH)2; and (bottom) no sample, onlyHCl. The letters denote steps in the experiment: A, container sealed; B,sample andHClmixed;C,maximumpressure; andD, container opened.
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described here is fairly simple, students have to recognize variousfactors that could influence its accuracy. This is useful as itprovides a deeper understanding of the problem than thatobtained from just calculations. There are several factors thatare difficult to account for when performing the experiment.However, some of these “sources of error” can be avoided with alittle effort. First of all, atmospheric humidity has a small yetnoticeable effect on the results as it increases the pressure insidethe glass jar. This source of error can largely be avoided byallowing for a long enough settling time before the stirring isstarted (Figure 2). Accounting for errors is also possible byrunning a few experiments without any carbonates and subtract-ing this result from all subsequent experiments. We have donethis and introduced a correction term called xcor (see thesupporting information). Another (small) source of error comesfrom pushing the rubber cap into the glass jar. Depending on theway it is inserted into the glass jar, the volume of the systemchanges (note that the problem can be altogether avoided if oneuses, for instance, a screw cap). However, the error is small andnot considered here in any more detail.
Discussion and Conclusions
Although there are various ways of determining the carbo-nation degree of a material, the method described here has thebenefit of being affordable and easy to use. In addition, itprovides a hands-on feel for the application of thermodynamicand chemical know-how to anyone who applies it. The fact thatthe method can be used for something of practical use works as amotivator in the otherwise often theoretical studies of chemicalengineers.
On the basis of our experience, students with basic knowledgein inorganic chemistry and some laboratory experience have noproblems adapting the methodology described here. The numberof instruments and computer loggers restricts the number ofstudents performing the experiment simultaneously, but perform-ing this experiment as a simple demonstration in front of theclassroom is also helpful. Still, perhaps the best results are achieved ifa group of students are given a task to determine the carbonate
content of a sample without any specific instructions or just a fewclues. For example, the students could be given the reaction in eq 1and left to work toward this method in their own way.
If the laboratory does not own a computer logger formonitoring the pressure and temperature in real time, themethod described by Pile et al. (7) could be used for determiningthe carbonation content, although we found their method lessaccurate with materials of low carbonate degree. A benefit oflogging the pressure and temperature data and plotting the datasimultaneously to the experiment is that it helps in detectingpossible errors immediately and predicting when the system hasreached a steady state. The method could be further improved byinvesting in a hygrometer that would fit together with thepressure and temperature sensor inside the glass jar. However,the method works well without this additional data also. Theaccuracy of the method is sufficiently high for determining thecarbonate content of mixed samples (e.g., MgCO3, Mg(OH)2,MgO) with a standard deviation of 1.1%.
This method functions as a good demonstration techniquefor students, providing insight into how the ideal gas law can beused and what assumptions are necessary when applying a theoryto practice in an educational environment.
Literature Cited1. Sipila, J.; Teir, S.; Zevenhoven, R. Carbon Dioxide Sequestration
by Mineral Carbonation: Literature Review Update 2005-2007;Report VT 2008-1; 2008; http://users.abo.fi/rzevenho/Mineral-CarbonationLiteratureReview05-07.pdf (accessed Sept 2010).
2. Huijgen,W. J. J.; Comans, R. N. J. Carbon Dioxide Sequestration byMineral Carbonation: Literature Review Update 2003-2004,Report ECN-C--05-022; 2005; pp 1-52; http://www.ecn.nl/docs/library/report/2005/c05022.pdf (accessed Sept 2010).
3. Williams, D. E. A RapidManometer Method for the Determinationof Carbonate in Soils. Soil. Sci. Soc. Am. Proc. 1948, 13, 127–129.
4. Choi, M.; Wong, P. Using a Datalogger To Determine First-OrderKinetics and Calcium Carbonate in Eggshells. J. Chem. Educ. 2004,81, 859–861.
5. Roser, C.;McCluskey, C. Pressure and Stoichiometry. J. Chem. Educ.1999, 76, 638–639.
6. Horvath, B.; Opara-Nadi, O.; Beese, F. A Simple Method forMeasuring Carbonate Content of Soils. Soil Sci. Soc. Am. J. 2005,69 (4), 1066–1068.
7. Pile, D.; Benjamin, A. S.; Lackner, K. S.; Wendt, C. H.; Butt, D. P. APrecise Method for Determining the CO2 Content of CarbonateMaterials. J. Chem. Educ. 1998, 75 (12), 1610–1614.
8. Elm, N.; Zipprian, J.; Schaber, K. Vapor-Liquid Equilibria of Binaryand Ternary Aqueous Systems with HCl, HBr, and CaCl2 at HighlyDiluted Vapor Phases. Fluid Phase Equilib. 2001, 189 (1-2), 163–178.
9. Ross, S. Peirce's Criterion for the Elimination of Suspect Experi-mental Data. J. Eng. Technol. 2003, Fall, 1–12.
Supporting Information Available
Details of calculating the carbonate content; how to correct forevaporating water. This material is available via the Internet at http://pubs.acs.org.
Figure 5. Known carbonation degree (from certified laboratory) versusmeasured results for five experiments.
Journal of Chemical Education 9/23/10 1 of 3
Supporting Information for Gasometric Determination of
CO2 Released from Carbonate Materials
Johan Fagerlund* and Ron Zevenhoven
Thermal and Flow Engineering Laboratory, Åbo Akademi University, Turku, Finland 5
