This lab was pretty simple, and was really just an excuse to burn more magnesium ribbon. We followed Ian Guch’s Energy Diagram Lab in 24 Lessons That Rocked the World. Students took an approximately 4cm long piece of magnesium ribbon, found its mass on the small pocket scale (it was too small for the triple beam balance) and then calculated how many moles of magnesium they had.
The lab says to put the ribbon in a crucible and place it over a bunsen burner. We don’t have bunsen burners so we put it over a butane burner and even opening it up all the way we could not get the magnesium ribbon to burn, so we ended up holding it with tongs directly in the flame. This very clearly showed the reaction had a high activation energy.
We can’t really measure the energy given off in the reaction but we can calculate it. The energy that would be released when they burned the ribbon is the amount of Mg they burned in moles (which was very small, like 0.004 moles) multiplied by the heat of reaction which is given in the lab in kJ/mole.
Students then sketched an energy diagram, showing that the products (MgO) have less energy than the reactants (Mg and O2) because its exothermic and it has a high activation energy (had to put it directly in the flame to ignite it). The lab had a worksheet at the end that the students worked on it in class since the lab was so short.
Students were asked to watch Crash Course Chemistry #32 Kinetics before class.
Continuing along in the Middle School Chemistry curriculum, today we did Chapter 6, Lesson 3: Forming a Precipitate. The lab starts out with a demonstration that forms a precipitate, which is a solid that forms from a chemical reaction. In this case, you mix some epson salts (magnesium sulfate) with water and some sodium carbonate with water in a separate cup. Stir both until the solute is dissolved then slowly pour the sodium
carbonate solution into the magnesium sulfate solution. A white solid forms, looking a lot like snow and it drifts down to the bottom of the cup (see photo above). In class we discussed how this was pretty clear sign that a chemical reaction had taken place.
For the lab the students got to make their own precipitate with calcium chloride (DampRid purchased from Lowes) and sodium bicarbonate (baking soda). They used a triple beam balance or the pocket scale to measure out 2 grams of each into small paper cups. Graduated cylinders were used to measure 20 ml of water into 2 clear plastic cups. All the cups were labeled with sharpies and the calcium chloride and baking soda were each poured into their own plastic cup and the students swirled them around until they were mostly dissolved. You can see there is still a little bit of baking soda in the bottom of that cup in the photo above, so when the students pour that cup into the calcium chloride cup they try to leave the residual baking soda behind. When they combined the liquids they observed bubbles and a white precipitate formed.
Students put the combined solution through a coffee filter to separate out the precipitate. The lab shows the students the chemical formula for this reaction and the products as calcium carbonate, sodium chloride, water and carbon dioxide. I asked the students which of these products was causing the bubbles and most guessed correctly that it was the carbon dioxide. Then I asked which product was the precipitate and that took a little more discussion because they know salt is a white powder but I reminded them that salt dissolves easily in water and there was no reason for it to precipatate out, which left the calcium carbonate (chalk!) as our precipitate.
This lab has the instructions for one more demonstration. Before class I made a copper II sulfate solution, which is a very pretty blue, and poured some of it into a test tube. To create a precipitate I used a disposable pipette to drop ammonia (10-20 drops), into the test tube. This formed a light blue precipitate which stayed on the top of the liquid and then a darker blue liquid layer formed above that. I also dropped in some hydrogen peroxide which made the dark blue layer even darker and occassionally brown, but I didn’t really see another precipitate which the lab said should form.
We had some time at the end of class so we watched these videos on youtube:
Crash Course Chemistry – this was way over their heads for some stuff but they asked to watch it.
We watched the first 10-15 minutes of the Royal Institute’s video on Fireworks. Its over an hour long but I figured if I got them hooked they would finish watching it at home.
This lab is pretty much a repeat of the lab last week but students had two different chemicals, potassium chloride (KCl) and calcium chloride (CaCl2), to dissolve in water while they observed the temperature change. Using the same equation as last time, Δ H = mc Δ T, where m is the mass of the water in the calorimeter (styrofoam cup), c is the specific heat capacity of water (4.184J/g°C)and ΔT is the temperature change, they calculated the heat of solution. The lab, IX-1: Determine Heat of Solution, in the homescientist chemistry manual used plain NaCl but a lab I did with the middle school class showed pretty big temperature changes with KCl and CaCl2 so I had the students use those instead and had them figure out how much solute and solvent (water) they should use for their experiment. There’s no real right answer for this, as long as the quantities they decide on are reasonable. The final number they’re trying to calculate is the molar heat of solution so if they use a tenth of a mole of the solute, they just have to move the decimal over one place at the end, which is what most of them decided to do.
For KCl, students used 7.5 grams of NoSalt, which is mostly KCl but not 100%, so the values we determined for the heat of solution are going to a bit different than accepted values for pure KCl. Same goes for the calcium chloride, it was not a pure chemical, but DampRid purchased at Lowes. The amount of water used varied from 50 ml to 200 ml, but all groups saw very noticeable temperature changes and got decent values for the molar heat of solution. I had each group do both chemicals because one is exothermic and one is endothermic.
The Calorimeter Lab is from Ian Guch’s 24 Lessons that Rocked the World and it involves building a calorimeter, bascially an insulated container, and measuring the heat of solvation for sodium hydroxide. I had styrofoam cups, newspapers, aluminum foil, bubble wrap and some insulated paper cups with lids out for the students to build their calorimeters. They had to make sure it could hold 200 ml of distilled water, a thermometer and that they could easily put in the sodium hydroxide (NaOH) pellets. Students measured the temperature of the water before dropping in the pellets (4 grams of NaOH) and then every 20 seconds until the temperature stabilized.
Students graphed their data, temperature as a function of time and calculated the heat of solvation (heat given off as sodium hydroxide dissolved in water) by using Δ H = mc Δ T, where m is the mass of the water in the calorimeter, c is the specific heat capacity of water (4.184J/g°C)and ΔT is the temperature change. We discussed this formula a bit before doing the lab. The heat energy that is released in dissolving NaOH goes into heating the water in the calorimeter and the more water you have the more energy you will need. The greater the temperature change, the greater the heat energy required. Its also easier to change the temperature of some things, like copper, while its difficult to change the temperature of water. So to calculate the heat required to change the temperature of something you need to know its specific heat capacity – the heat energy required to raise the temperature of 1 gram of that substance by 1°C.
For this experiment we expected to get results that were a little below the predicted value, showing that we actually lost heat to the surroundings and our calorimeters weren’t perfect. But everyone ended up with results on the high side and we couldn’t figure out what went wrong. We suspected that the pocket scale we used to measure the 4 grams of NaOH pellets might have malfunctioned, but checking it with known masses and a triple beam balance it appeared to be working correctly. Everyone confirmed that they used the correct amount of water, filling two 100 ml graduated cylinders with distilled water. We calibrated our thermometers a few weeks ago so I don’t think they were the problem. I would have liked to repeat the lab, but we used up all our sodium hydroxide – which is another reason I suspected our scale because we should have had some left over.
The only dangers with this lab are the NaOH pellets and making sure to neutralize the NaOH solution with vinegar at the end of the lab before dumping down the sink. I used some left over cabbage juice indicator so I would know when the solution was neutral. Instructions for this are included in the lab handout.
Students watched the following Crash Course Chemistry videos before class:
Today’s lab, Controlling the Amount of Products in a Chemical Reaction, is lesson 2 from Chapter 6 of the ACS Middle School Chemistry Curriculum. All you need is a graduated cylinder, some measuring spoons (or scale),vinegar, dish detergent and baking soda. At the beginning of the class I put 1/2 teaspoon of baking soda into a 50ml graduated cylinder and then pour in 10 ml (measured earlier) vinegar with 1 drop of detergent solution (1 tsp of detergent + 2 Tbsp water). If you’ve ever done the volcano labs you know this is going to react and make lots of bubbles/foam and indeed it bubbles up and out of the graduated cylinder. The detergent is just there to help form the bubbles/foam and is not part of the chemical reaction. The handouts I gave the students had the chemical reaction on it: acetic acid (vinegar) + sodium bicarbonate (baking soda) produces water, carbon dioxide and sodium acetate. The challenge I gave them was to find just the right combination of vinegar and baking soda to produce enough bubbles (carbon dioxide) to come just to the top of the graduated cylinder without overflowing. A couple of groups managed it quite quickly so I then gave them a 100 ml graduated cylinder and asked them to repeat the experiment for that cylinder. One girl immediately realized it was twice the volume so doubled their previous results and got it the first try. The handout says to use measuring spoons but I didn’t have any smaller than 1/2 tsp so I had the kids measure the mass of the baking soda in little paper cups.
For groups that finished early I had snatoms and zometools out so they could build the molecules involved in the reaction. After everyone was done I asked if we could just keep making larger quantities of carbon dioxide (bubbles) by just adding baking soda to the 10 ml of vinegar. Most of the students realized that no, you would need to add both baking soda and vinegar. I told them it was like when you go to make cookies and find you only have 1 egg and the recipe calls for 2 eggs. You can cut your recipe in half and make half a batch of cookies, but you can’t just use 1 egg and the same quantity of all the other ingredients. And if you want to make two batches of cookies, you’l need to double all the ingredients (and go buy more eggs).
George Lakoff has retired as Distinguished Professor of Cognitive Science and Linguistics at the University of California at Berkeley. He is now Director of the Center for the Neural Mind & Society (cnms.berkeley.edu).