# Tagged: inquiry

# Playing with Chemicals

This week, I had the AP students do some very simple labs focused upon law of conservation of mass, but connecting it to stoichiometry. Students mixed a white powder with a liquid to produce a gaseous product (massing beakers initially, with and after, see image at right), share their data on the board, then graph the class data on a TI calculator (see image below).

There were a couple of points to this. First, the initial share was with total mass before and after and the students see they don’t match — the gas was *lost.* This means different data is needed, but it still has to come from what they collected. Why not compare the mass of the gas (that *lost* mass) and the mass of the solid? When this is graphed (solid mass on the x-axis, gas on the y), a linear relationship is evident; further supported by the linear regression done. [I’m reading my notes from this summer, one-finger-typing my lists and such and my students were waiting on me; even made a point of telling me to just ask *them* what to do next time as they do this all the time in math. They enjoyed that way more than they should have.] We do discuss just what the linear equation means (y=mx+b). In this case, y = 0.502x +.055, the slope tells them that 0.502g of gas is produced for every 1g of solid used and that there’s error because ‘b’ has a value other than zero (the error in the class data).

Then, the second point of the discussion — the part I did not carry off as well as I wanted — determining the ‘ideal’ relationship for this solid and this gas. Enter the chemical reaction (just in case they had not determined it was baking soda and vinegar) and I place the molar masses of each substance underneath, totaling the masses of reactants and products to show the masses are equal, show conservation of mass. This wasn’t exactly a revelation to this group, but I didn’t **let them tell me** what the ideal equation for the line should be: the total mass of the gas divided by the total mass of the solid. This is the point that gets us to stoichiometry. As soon as you start comparing the masses, ideally seeing that molar ratio behind the mass comparison because *molar masses were multiplied by coefficients*, you are performing the basics of stoichiometry calculations.

They answer four questions

- How good is your point? Explain. [Their point on their graph has a smiley face above it.]
- How many grams of gas should you lose? Explain how you know.
- ? g of gas if 5g solid used? Show 2 ways.
- ? g of solid if 5g gas produced? Show 2 ways.

All of these are answered by evaluating the graph and their point on the graph

- it’s good because it’s close to the ideal line
- more or less gas should have been lost — based on whether it’s above or below the line (getting to another rabbit trail in the discussion about the source of error: below the line, not enough gas produced; above the line too much; and what could have caused this)
- the first way, find it on the graph; second way, calculate by multiplying masses (factor label)
- the first way, find it on the graph; second way, calculate by multiplying the masses, but inversely as the gas needs ‘canceled’

The goal with this set of labs is to have the students (1) want this ‘proved’ again to see this truth with a different reaction; (2) lead to other questions like what the graph would look like if the amount of solid is varied with the same amount of liquid or vice versa, keep the solid constant and vary the liquid; i.e., predicting limiting reagents. Again, the whole point is to get them to take the data, see the linear relationship, but ask why and want to see more data to prove it — not exactly carried off in this way. Instead, I had them do the lab several times, plug in our data, examine the graphs produced, and see that, “yup, it’s linear, again.”

My problem here was my approach. This would work well for chemistry if I fix the conservation of mass introduction & link to the results — what does conservation of mass mean? reactants should equal products, so we balance the equation, but is there another way?…work to masses of reactants and products and what this means for the lab, compare the line produced to the ideal one *they* find…ask if this is the case all the time? is it predictable? Push for the questions because it’s new. AP, on the other hand, should be about predicting the results, predicting the graph, predicting the ideal line. This information is not new it’s review, so, I needed to take it to that point instead of getting caught in no-man’s-land between new & review where I lost some of the effectiveness of the enterprise.

[Thanks to Jim Cortez for sharing this during our summer APSI.]

# Exploring Chemistry

I’m liking the challenge and the potential impact of ‘exploring’ each concept before students actively study as I attempt to implement explore-flip-apply. My twist, though this is blatantly stolen, is to have the students reflect on &/or explain the exploration in a blog post.

My post is coming up prior to any data to support an assertion of actual improved outcomes and increased learning. However, the questions asked, exposing the interconnections between concepts, was absolutely amazing today.

The activity, using a 9-volt battery to eletrolyze water, is part of my first unit and I wanted it to both review and set the stage for stoichiometry. There were four questions I wanted them to answer:

- What is the balanced chemical equation?
- Is there qualitative evidence to support the balanced chemical reaction?
- Could you collect quantitative data to ‘prove’ the balanced reaction? How?
- Can you draw a particle diagram(s) that models what’s going on in this reaction?

I had the equipment out — battery, sample cups with tacks, small plastic test tubes, and two different salt solutions — and we got started. One of the first questions asked was how to capture the gas in the test tubes. This is not a question to be taken lightly, since the point was to have the captured gas push water out of the test tubes to visually see the difference in the amounts. Rather than let the students struggle, I made a mistake I think, I showed them what to do — fill the test tubes with the salt solution, invert, and quickly fill the sample container with more solution. The whole apparatus is now placed upon the battery. Immediately, bubbles begin forming, an unmistakable difference in rate apparent. The students get theirs going.

I wanted them to work alone to answer each question first, thinking about them while they watched the reaction, then using their ideas during discussion. Again, I think I jumped the gun a bit — struggling is not something they enjoyed — and I cut this time too short.

We jumped into the group discussion with the first question and a uniform response was provided, a balanced equation for decomposition of water. I jumped to the last question here, I’m leaning toward making it the second question next time, and, again, a confident reply of ‘sure’ from the group. The second question was the first divergence from my script: how can you know that the gases are actually hydrogen and oxygen? The observations also helped to push this question forward from left-field. After a bit, all the test tubes lost the apparent doubling of gas in one test tube versus the other; there was still more in one, but it didn’t look like twice as much. To try to show this, I introduced some UI to the solution and filled the tubes and sample cup again. One complication, the salt solution used sodium bicarbonate.

Shifting our focus again, during this time of waiting and watching, we jumped to the third question; surprisingly tougher than I thought. They were still focused upon how to measure the products, how to verify they were oxygen and hydrogen….I just wanted them to think how a balanced equation had to be based on an equal mass before and after. So, I kept trying to push them back to the law of conservation of mass, the law behind a ‘balanced chemical’ reaction. Again, I gave in, and just told them this.

Concluding the UI variation….

Gases are now being produced in a blue solution, but bless it, one of them begins to lighten (I won’t go so far as to say it turned yellow, but it did become less blue). This lead to what exactly this change in color meant. We take a turn into pH, the equilibrium of water ionization and baby steps to electrochemistry. Using the equilibrium equation as a new starting point, I try to encourage them to work out what’s ‘left’ when each gas is formed, pushing them to visually separate the equation in their mind and the bell rings.

I try to frantically throw information at them as they ready for the next class and assure them we’ll do a quick finish-up tomorrow, in class, and their blog post will be due tomorrow now, too.

[Follow-up: so quickly refocused upon goals from yesterday, added some explanation about self-ionization of water, rewriting the equilibrium equation twice. Without going into detail on redox, so just in terms of particles, if hydrogen is removed (or oxygen), seeing what is left behind helps to explain why the indicator changed color, why the pH is different. I’m really hoping this turns into a seed to reap from in future concepts.]