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R2 Physics 15 – Ideal Gas Law

The only new physics ‘toys’ I bought for teaching this year were some Go Direct sensors by Vernier.  They work directly with an iPad via bluetooth and their cost was fairly reasonable ($59 & $79 for the ones used in this lab) for a small class.   For today’s lab the students needed to measure the pressure of a gas as they changed its volume (at constant temperature) and as the temperature changed (with constant volume).

The first set up is very simple, as seen in the photo below, its just a syringe attached to the Go Direct pressure sensor.  The syringe actually screws and locks onto the sensor.  Students started with 10 ml of air at atmospheric pressure and room temperature in the

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syringe and then attached it to the sensor.  One student in each group connected their iPad to the sensor via bluetooth and the Graphical Analysis app.  They then changed the volume of the gas sample by pushing the plunger in on the syringe and recording the pressure and volume.  They also took data at lower pressures by increasing the volume, pulling the syringe out.  The volume is easily read off the syringe.

P vs V graph

The data came out great.  Some students graphed the results by hand, some used the Data Analysis app or the Graphical Analysis app, did a fit to the data and found that pressure did indeed depend on 1/Volume. Solving the ideal gas law for P (pressure), you get P = nRT/V and in this experiment nRT are all constant (n = number of moles, R = ideal gas constant, T = temperature).

Since I only had one pressure sensor the students had to wait and take turns doing the experiment, but it only took about 10 minutes to take the data you see in the graph on the right.  I put the next homework assignment on the board and had students work on that until it was their turn to take data.

This worked sooooo much better then when we did this lab in the past using a syringe and balancing books on it, calculating the pressure on the syringe from the weight of the books and then measuring the volume. It just never came out very well so I’m very glad I bought this sensor, it made this lab a lot easier and it actually worked!  This is one of the few labs you end up doing in both chemistry and physics class so I should get plenty of use out the pressure sensor.

For the second lab, we used two Go Direct sensors, the pressure sensor and a Go Direct temperature sensor.  We could have just used a thermometer, but I bought the temperature sensor to use with some other labs we’l be doing in the spring and figured I might as well use it. I’m pretty sure the iPads will only connect to one sensor at a

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time so we had one student measuring the pressure and another student measuring the temperature.  The pressure sensor came with the stopper and tubing seen in the photo to the left.  It fit a flask I had and that became our gas sample for the second experiment. We did this experiment as one big group.  The first data point was at room temperature, then we placed the flask in an ice water bath to get a lower temperature data point.  Then we put the flask in a warm water bath and a hot water bath for a total of 4 data points. You can see the results in the graph below,  pressure depends linearly on the temperature of the gas when the volume is held constant. From the fit you can also see that the y-intercept, b, is pretty much zero, which is what you would expect. Pressure should be zero at T= 0 K, absolute zero.  P vs T graph

From the ideal gas law, P = nRT/V, in this experiment n (number of moles) and V, the volume were held constant and we see pressure equals a constant times temperature.  I’m very happy with these sensors, this lab was always a chore and now its so easy.  I’m going to have the students write up the results of this lab in a formal lab report.

At the end of class I brought out a windbag  – basically a 2 meter long thin bag, roughly 6 inches in diameter.  I laid it out on the table and asked my students how many breaths they thought it would take for me to blow it up.  Most immediately said a LOT of breaths, but one already knew the trick and said I could fill it with one breath – which is the right answer!  By holding the end of the bag wide open and having the bag laid out on the table, I had my mouth about a foot from the opening of the bag and blew into it.  Blowing into the bag at a distance lowered the pressure around the mouth of the bag causing the air around it to rush into the bag – hence filling it with one breath.  I had a box of these bags and each student got to take one home. These were more fun than I expected and a nice demonstration of Bernouilli’s principle. IMG_4996

 

 

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Honors Chemistry 32 – Specific Heat

IMG_0123We managed to plow through this course in pretty good time and the only chapters left in the text are organic chemistry and  since most of the class already did biology with me two years I decided to skip those chapters.  Since we managed to get through all the labs I had planned, I went back and decided to do another specific heat lab from the Home Scientist Chemistry kit manual CK01A, Session IX-3: Determine the Specific Heat of a Metal.  This lab suggests using 25-50 pennies, find their mass and then place them in a beaker with roughly 100 ml of water.  Bring the water to a boil for a few minutes and then measure the temperature (of the hot water and hot pennies).  IMG_0124Remove the pennies from the hot water and place them in a calorimeter (styrofoam cup) containing 100 ml of room temperature water (measure T before putting in hot pennies).  Measure the temperature every 30 seconds until it stops going up.  The heat gained by the room temperature water is equal to the specific heat of water (4.184 J/gK) times the mass of the water (100g) times the change in temperature (roughly 4 or 5 degrees).  Since the heat gained by the water is equal to the heat lost by the pennies we can use the same equation to solve for the specific heat of the pennies.  The students got fairly good results for this lab and it was pretty quick, only took an hour or so.   I happened to have a set of metal cylinders for density labs and some of the students used those instead of the pennies.

That’s it for this class.  I hope you found these posts useful and if you have any questions feel free to post in the comments.  IMG_0122

Honors Chemistry 31 – Nuclear Chemistry

This is one of my favorite chemistry topics since its also a topic in physics.  I pulled out an old slide show on nuclear physics, different types of nuclear reactions, fission, fusion, atomic bombs, power plants, all the good stuff.  We did the activities ‘Simulation of Nuclear Decay Using Pennies and Paper’, from the Modern Chemistry curriculum and built cloud chambers.

For the paper activity,  I precut a bunch of paper strips from colored card stock and gave each IMG_9880student two strips.  They placed the first strip on a graph to represent 100 percent of material.  Half of that will decay in one half life so they take the second strip which is the same length as the first one and cut it in half.  Tape the half strip next to the first one.  Repeat with each remaining strip until you can no longer easily cut the strip.  For this example we made the half life 1 minute, which is about the time it took to fold and cut the piece of paper to make the next bar.

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The other part of this activity involved putting 100 pennies heads up in a box – this represents our ‘original’ sample of material.  Students shook the covered box 5 times and then removed all the pennies that had ‘decayed’ (turned to tails).  This should be roughly half the pennies.  They then repeated the shaking and removing of ‘decayed’ pennies til they had 1 or 0 pennies left.  Every shake was considered to be a half life of 10 minutes for the purposes of graphing their data.

IMG_9892Then we get to the fun part of class, making cloud chambers.  A cloud chamber is a closed container with an isopropyl alcohol soaked felt pad inside it (near the top or sides) and black paper on the bottom.  The alcohol forms a mist inside the container becoming super saturated near the cold bottom of the container which is sitting on dry ice.  As particles zip through the mist it produces ions and the alcohol drops condense on these ions, leaving a visible trail.  There are a lot of videos on the web explaining how to make cloud chambers.  Here are the two I like, one by Jefferson Labs uses a petri dish to make a small cloud chamber and the one on ScienceFriday has instructions for using something bigger.  The petri dish one works really well but its gets fogged up and you spend a lot of time wiping it off.  I bought the dry ice and 91% isopropyl alcohol at my local grocery store. Our best cloud chamber was built from a cheap (thin and flimsy) plastic cookie (Dunkers) container from Trader Joe’s.  IMG_9917

In the photo above you can see our radioactive rock and 4 trails from particles that were emitted from the rock. You can also see the alcohol mist/rain in the container.   You don’t need a radioactive sample to put in the container, you will see trails from muons and other particles that are zipping by us all the time.   Here’s two videos from class:

I asked students to watch these videos before class:

Honors Chemistry 30 – Review lab

We’re almost done with this class and actually finished all the labs that I had previously pulled aside so I spent some time last week browsing the labs that come with the Modern Chemistry homeschool curriculum  and found a few more to try.  Today, we did “How much calcium carbonate is in an eggshell?” from Chapter 15,  Modern Chemistry.  We didn’t get the expected result but its a really good lab and uses a number of concepts including stochiometry and titration.

To start you need a clean eggshell with membranes removed.  Bake in an oven at 110F for 15 minutes – my oven can’t be set this low so  I set it to 160F then turned it off and put the eggshells in the oven for 10 minutes.  Once the eggshells are cool, grind them up to increase surface area.  Put 0.1 grams of powdered eggshell in a small beaker (50ml) and then add 6.0 ml of 1.0M HCl.  The lab handout actually has students figure out the volume of one drop from a pipette and put in 150 drops of HCl into the beaker but since we had pipettes marked in 1/2 ml increments I decided to do it the less monotonous way and just put in 6.0ml.   Students swirled the beaker and watch the reaction.   After 4 minutes, two drops of phenolphthalein indicator were put in the beaker.  Phenolphthalein remains clear for acids and neutral solutions, but turns bright pink in basic solutions.

IMG_9770Students then added 1.0M NaOH solution in 0.5 ml increments into the beaker to neutralize the remaining HCl that did not react with the eggshell.  Both groups ended up putting in 6 ml of the NaOH solution which meant that none of the HCl reacted with the egg shell, which couldn’t be right since we saw a reaction take place (bubbles!).  So we tried again but this time we used red cabbage indicator, which has a range of colors depending on the pH.  The students also used 10ml graduated cylinders to measure the HCl they put in initially and started with 6.0ml of 1.0 M NaOH in a graduated cylinder and slowly moved it to the beaker with a pipette. This way we could just measure the remaining NaOH and know how  much we had put in the beaker.  Unfortuanately we still ended up putting in all the NaOH, only with the last few drops did the indicator record a significant shift in pH.

IMG_9784So then we thought about the titration process and how it was only going to work if the the solutions were both 1.0 M and perhaps we had done that wrong, so we started from scratch and remade the 1.0M solutions from the 6.0M HCl and 6.0M NaOH (These chemicals come in the Home Scientist Chemistry kit)  This time we used the same graduate cylinder to make both solutions.  We used the cabbage indicator and very carefully swirled between drops for the last bit NaOH and found that yet again we had to use almost all the NaOH, 5.6 ml, which means only 0.4ml of HCl reacted.  We went ahead with the calculations and found that our eggshell was only about 20 percent calcium carbonate and the expected number is closer to 80 percent.  I’m not sure what went wrong – did I over heat the eggshell, breaking down the calcium carbonate?  Don’t think so since it really didn’t get very hot.  The eggshells were sitting out on a counter for a few days before we did the experiment (and before heating) so next week I’l have the students  try fresh eggshells.  Perhaps our 6.0M solutions have been contaminated or aren’t quite 6.0M?  Were the eggshells not crushed finely enough?  I’m not sure what went wrong but since this lab didn’t take very long to perform we were able to do it three times and tried to improve the procedure each time.  Blair Lee from SEA Homeschoolers just posted last week about ‘When Experiments Don’t Work, That’s When the Science Really Gets Fun!’, which is exactly what happened today.

Honors Chemistry 29 – Electrochemistry

The labs we did today came from the Home Scientist CKO1A instruction manual which goes along with their chemistry kit.  In Session X-1: Observe Electrolysis there are two labs, in the first one you set up two test tubes full of water (with some epsom salts) upside down in a beaker and then place the wire ends of the battery adapter into the tubes.  Once you connect a battery, current starts to flow through the water breaking it down to create hydrogen and oxygen gases in the test tubes.  We tried to set it up like the lab describes but the battery adapters I bought had short leads and we got pretty frustrated with it.  Luckily I happened to have two electrolysis set ups (see photo below)  that were loaned to us by another homeschool mom.  IMG_9692

The black stand in the bottom of the beaker holds the test tubes in place and the screws are the electrodes, so all we had to do was connect the wires to the battery (after filling the tubes with water).  The other advantage of using this set up is that the test tubes were very small so it didn’t take as long to fill up with gas (though it still took almost an hour).    We skipped measuring the volume of the gas but did notice that one test tube filled up twice as fast as the other because you make 2 hydrogen molecules for every oxygen molecule produced.  The blue arrows in the photo above point to the water level in the test tubes.

While this was bubbling away, I set up the second part of the lab which involved doing electrolysis with salt water.  We took 50 ml of water and a tsp of salt, stirred until it was dissolved and then made electrodes out of Al foil.  We took this lab outside since it produces chlorine gas.  We also put a few drops of phenolphthalein pH indicator in the salt water.  Phenolphthalein is clear in neutral solutions but turns a bright pink in basic solutions.  I hadn’t done this before so I was as suprised as the students when I hooked up the battery and the solution turned pink starting at one electrode and making its way away across the beaker.  The bubbling gas production was also quite vigorous.  We weren’t prepared for how quickly this took place so I had the students repeat it themselves and we took lots of movies and pictures.  If you leave it hooked up for more than a few seconds the Al foil starts to break apart.   When the students did it, the reaction didn’t seem quite as vigorous and that’s because I used a regular teaspoon to put in the salt and they used the chemical spatula which gave them less salt.  Here’s a video of the electrolysis of salt water.

The solution turns pink because sodium hydroxide (a base) is formed along with hydrogen gas at the foil electrode (cathode)  where you first see the solution turn pink.  Chlorine gas is formed at the anode but most of it dissolves into the water.  Very cool little experiment.

I had the students watch Tyler Dewitt’s video on Electrochemistry before class.

He also has this great video which explains exactly what was happening in both labs we did.

 

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