Friday, February 10, 2017

Gravitational analysis: AP Chem vs. BU Analytical Chem

Chalk!
Here is another blog post from my summer Analytical Chemistry class at Boston University. I dusted off this essay just in time for my students to tackle gravitational analysis in the lab this week.

During the summer course I had the chance to do the “college” version of the classic gravitational analysis lab that I like to do with my AP Chemistry class. This experiment shines a spotlight on the net ionic equation, in this case calcium ions reacting with carbonate ions in solution to form insoluble calcium carbonate.  I love that my students can see the insoluble product “falling out” of solution (it happens in a massive dump in the high school version). How better to give solubility rules meaning than to see two clear solutions mix to form a white precipitate? But the other powerful aspect of this double displacement reaction is the stoichiometry that reveals the identity of an unknown substance. The mole ratio in the chemical reaction takes on tangible significance in this lab application.

Vacuum filtration is a new addition for me at Woodstock Academy.
The AP version of the lab I like to use with my students is called “Gravitational Analysis of a Metal Carbonate.” (Laboratory Experiments for Advanced Placement Chemistry by Sally Ann Vonderbrink) In the lab, the students choose an unknown alkali carbonate (potassium carbonate or sodium carbonate, I tried lithium carbonate but it doesn’t dissolve easily). The students dissolve a known mass of their metal carbonate in water, and then precipitate out the carbonate ion with an excess of calcium chloride solution. The reaction between the calcium and carbonate ions happens in a mass precipitation that occurs instantly to produce a very fine powdery calcium carbonate product (commonly known as chalk). Even with electronic balances that only measure to a hundredth of a gram, the students can collect sufficient data to identify their alkali metal carbonate. However, this year at Woodstock Academy we will be using our wonderful analytical balances that measure to the thousandths place! By measuring the mass of the product formed, after it has been collected and thoroughly dried, the students can work backwards from the moles of calcium carbonate product to the moles of carbonate ion and ultimately to the moles of metal carbonate in the original mass. By taking the ratio of mass to moles of the original substance, the students can determine the molar mass of their unknown and ultimately the identity of the alkali metal carbonate.

The calcium ion solution with unknown concentration.
Kicking this lab up a notch to a college analytical lab level was a lot of fun. So what’s different from my high school AP Chemistry and the college version? Well, everything and nothing, all at the same time. 

We started with a calcium solution that contained an unknown concentration of calcium ions (well, the lab technician knew the calcium ion concentration, but she wouldn’t tell anyone). Our goal for the lab was two-fold: figure out how much calcium was in the solution, and do statistical analysis of the results to decide if our answer was worth the “paper” it was printed on.

The urea was used to create a slower reaction, and the methyl red indicator told us when it was done.
In this version of the experiment, we had the time to do a more sophisticated precipitation reaction that used both particle growth and homogeneous precipitation to make larger insoluble calcium salt particles. We set up a clever reaction of calcium with a “time released” oxalate ion to form the insoluble calcium oxalate precipitate. By producing the oxalate ion in a slow and steady environment from the reaction of ammonia with oxalic acid, the precipitation reaction happened relatively slowly.
We heated our reaction mixture on these giant hot plates for about an hour.
We cooked our reaction mixture for about an hour on a hot plate in the fume hood. We knew the reaction was done when the methyl red indicator in the reaction mixture changed from red to yellow. Methyl red is a pH indicator that gives a visual signal that the reaction between the urea and the oxalic acid has reached completion. Methyl red has a red color when it is below pH of 4, and it turns yellow above a pH of 6. When the pH of the solution gets high enough to see the yellow color, then the oxalic acid has been fully consumed and the reaction is complete.
The yellow color appeared when all the oxalic acid was consumed by the reaction. Time to filter!
This reaction gives a slow and steady production of the oxalate ion, allowing the precipitate to form at a slower rate compared to the massive dump in my high school lab. Although the product was the same white powdery product that I see in my AP Chemistry lab, the larger particles were easy to collect using vacuum filtration. We made up three identical reaction mixtures and collected three separate calcium oxalate products in our vacuum filtration. In the end, we were the proud parents of calcium carbonate triplets.
Gotta love the vacuum line at the bench.

Water is the enemy in this experiment because it can add unnecessary mass to all the sample measurements, giving inaccurate results for the calcium ion concentration. To get everything good and dry, we used pre-heated funnels for our vacuum filtration, along with a series of drying steps to thoroughly dry our product. The cycle went something like this: heat the dickens out of the samples in the microwave for 5 minutes, cool for 15 (measure the mass), heat in a drying oven for 5 more minutes, cool again (measure mass), repeat until three consecutive mass measurements were the same.
Here are the dried solids on our filters. Water is the enemy in this phase of the experiment.
There is a reason why these labs are four hours long! I used all the drying cycle wait times to write the answers to the post-lab assignment questions, with fill in the blanks for sentences with data and results. I’m sure that none of these undergraduate student in my class have three kids at home waiting to go to the swimming pool as soon as they get out of the car. It took three twenty-five-minute heating cycles to get my samples dry enough to record a final mass measurement. In my high school lab, we leave the precipitate on the filter paper in the drying oven over night (sometimes for several nights if it’s a Friday) to get the solid good and dry.

Recording accurate mass measurements was critical in this experiment. To do the job right, we used high precision analytical balances housed in the “balance room,” a stuffy room with no ventilation in the back of the lab. In this room stood about a dozen analytical balances, just waiting for us to use. I really loved using the analytical balances in the balance room.
The balance room, where students make very accurate mass measurements and plans for the weekend.
Not only were the balances excellent lab equipment at our fingertips, but also the balance room was one of the only places that the students interacted during lab. Most of the four hours were spent working at the bench in silence. I never expected undergraduates to be so shy! After all the drying and massing, and massing and drying, we recorded our final mass of the precipitates. And then, after one last nice look at our white salts, we unceremoniously discarded our calcium carbonate triplets in the trash.

The second objective in our analytical analysis of calcium solution was met the old fashioned way: statistics. We wrote our three final masses on the chalk board (public humiliation happens in college labs too!) to consider the quality of our results. I was surprised by the wide range of masses that we collected from the same chemical reaction, and I mean identical reaction conditions from start to finish, right down to measuring out equal volumes of the unknown calcium solution with a volumetric pipette. I wondered how this could be possible. When I looked around the lab, all of my classmates were doing the same steps with the same equipment; they looked competent and seemed to understand the procedure. I was baffled by the wide range of outcomes. Katherine assigned everyone two additional sets of data to use to do the Grubbs test. Using the Grubbs test, you can compare data points with the average and the standard deviation of the data set. Using this simple test, you can decide if a data point is statistically valid. Any points that “failed” the Grubbs test would be discarded from the data set to get a better final answer for the calcium ion concentration. We were purposely assigned data sets that were out of range so that the calculations would show something. Why not make it a teachable moment? With the same spreadsheet we could have crunched Grubbs test for the data from the entire class, I’m not sure why we didn’t. I always tend to lean towards more data, but I kept my mouth shut on this idea so as not to get throttled by my undergraduate classmates, who were ready to get out of the lab. My data was pretty consistent with the two other sets of data Katherine assigned me to analyze. None of the data points were discarded based on the Grubbs test which made my analysis a bit easier.

Table 3.3 – Average concentration of Ca2+ in unknown solution (g/L) and the calculated Grubbs test value.

Trial
Conc Ca2+ (g/L)
Grubbscalc
1
6.534
0.388
2
6.538
0.035
3
6.517
0.009
4
6.188
0.041
5
6.428
0.005
6
6.087
0.060
7
6.589
0.020
8
6.659
0.033
9
6.585
0.019


The statistical analysis of the data set was an important lesson for these students who are planning to go on to study advanced science and possibly pursue a career in scientific research.  Dr. Abrams made a good point to the class in our pre-lab lecture about the importance of what I call scientific checks and balances. When these students go out into the field to collect original data, there will not be a standard for comparison. In “the real world” a scientist must use these statistical tools to validate their results and communicate to others that their methods and conclusions have merit. This point may have been lost on my young classmates, but as a science teacher this argument really stuck with me as a good reminder of why we do scientific research and how we decide if information is meaningful.

Taking this deeper dive into gravitational analysis only solidified my love for the high school version of the experiment. I am excited for my students to get the chance to “reinvent” this experiment when they take it at the college level. Hopefully they will remember more about the experiment than waiting around for the white powder to settle. 

Monday, January 2, 2017

December NEACT Meeting: a Fresh Look at the Atomic Model

Teachers from CT, MA, VT, and RI came together today to discuss/explore the atomic model.
Twenty teachers from around New England gathered at the Woodstock Academy chemistry lab for a morning of atomic model discussion. Emily Allen Hoffman from Boston University kicked off the program with insights into student misconceptions concerning the atomic model from her research. Emily’s presentation was an excellent reminder for all of us just how difficult it is to teach the atomic model. Over the course of two years Emily interviewed Boston University chemistry students about their understanding of the atomic model. Her research showed that students who have exposure to a wide range of visual and graphical models to depict the atom are better able to move past their misconceptions and communicate effectively their understanding of the atom.
Emily Allen Hoffman from BU presented her findings about students' misconceptions.
As part of her doctoral work, Emily developed a series of lessons for the BU chemistry students to use concurrently with their regular course work. The students who had access to these interactive lessons, which include interactive models and demonstrations, were more successful in the course and gained a more complete understanding of the atomic model.
Emily asked the teachers to draw an atom.
The teachers at the workshop engaged in a lively discussion about the challenges we all encounter when teaching this abstract concept to students. We all agreed that the misconceptions Emily uncovered in college students were present in our classrooms in high school and middle school science. Armed with more ideas and resources, we all got a boost of energy for tackling the atomic model with our students.
We all need to remind ourselves how hard it is to teach the atomic model.
In the second half of the morning, Carline Chute and Mel Gronski from Woodstock Academy shared several hands-on activities that they do with their students. The activities included "atomic bowling", building models with Legos, Build an Atom Phet simulation, and an atomic model web-quest. After a demonstration of each activity from Caroline and Mel, the teachers had the chance to try out the activities and discuss how to adapt them to their own classrooms. Every participant left the workshop with ideas and activities to implement immediately in their own classroom.
Caroline explaining how she uses physical objects like legos to help students visualize the atomic model

Mel is explaining how she uses a WebQuest with her students

Atomic bowling is a fun analogy for the Rutherford gold foil experiment.

These teachers are exploring the PhET app on the iPad.

Teachers from all phases of their career came together today to discuss teaching chemistry.

The best part about NEACT meeting is meeting new chemistry teachers.


Trying out the "Build an Atom" PhET simulation for iPad.
Emily and Caroline are two of my favorite people, what a joy to have them in the same room for the morning!

Another look at the group of enthusiastic chemistry teachers at the Dec. 3 NEACT meeting at Woodstock Academy.