High School SMILE Meeting 1999-00 -- 05-06 Academic Years Properties of Gases

09 November 1999: Marva Anyanwu (Green School)
phenomenologically involved us with air pressure (handouts). At one point, Ken Schug (IIT) lifted up one end of a desk with a suction cup (dent puller). And we had an interesting discussion of physiological effects of changes in atmospheric pressure. And - using drinking straws and ziplock bags we raised books up. Insert the straw into the bag, sealing with tape. Then blow into the straw, inflating the bag - which is under the book and is so raised up. Good ideas!

27 February 2001: Pat Riley (Lincoln Park HS) Gas Laws
Non-Mathematical Part:

• She placed a plunger of a syringe on a convenient and simple setting, say 10 or 20, corresponding to a particular volume.  We predicted what would happen to the volume if we held our thumb over the opening and pushed on the plunger.  There was general agreement that the volume would decrease.
• Boyle's Law tells us that, as the volume of fixed amount of gas decreases, the pressure increases---where pressure is defined as  force per unit area
• Now we placed a hand on a piece of paper and traced its outline.  When we push on our hands, we are applying pressure only over the region traced on the paper.
Part involving some mathematics:
• Set a large syringe set at a volume of  100 cm3. Then place the tip onto a rubber stopper to make a seal.  We applied additional pressure by placing books on the stopper, measured the resulting volumes, and made the following table:
 Pressure (Number of Books) Volume (cm3) Inverse Volume (cm-3) 0 100 10.0 ´ 10-3 1 97 10.3 ´ 10-3 2 83 10.8 ´ 10-3 3 86 11.6 ´ 10-3 4 75 13.3 ´ 10-3 5 70 14.3 ´ 10-3
• Kinetic theory states that all gaseous matter is composed of particles that move randomly in straight paths until they collide.  They then bounce off in a new direction, until the next collision.  Collisions of the particles with the walls of the container produce pressure.  As you reduce the volume, there are more collisions with the wall, and hence more pressure.
• As you increase the pressure on the syringe, the gas particles become closer and closer together.

More mathematical part:

To estimate the absolute pressure of the air in the room, make a graph of  Pressure (Number of Books) versus the Inverse Volume (cm-3). The results will look like this
```
6  |
Pressure   |                                  *
# books    |                                *
3  |                              *
|                            *
|                          *
|________________________*_________
|    2    4    6    8   10   12   14  ´ 10-3
|                    .
|    1/V           .
-3 |                .
|              .
|            .
-6 |          .
|        .
|      .
-9 |    .
|  .
|.
-12 |

```
Atmospheric pressure [as measured in "number of books" on syringe] corresponds to the crossing point on the vertical axis; the vertical intercept,  or about 11 or 12 books.

Next presentation: Charles Law

01 May 2001: Sarah Brennan (Robeson HS)  Handout:  Properties Common to all Gases
Sarah gave 35 cc syringes and balloons to each of us.

• Gases have mass:  We used 2 balloons, one inflated with air and tied them with string to form each side of a balance.  From the fact that that balance tilted toward the inflated balloon, we concluded that the inflated balloon has more mass, because of the air inside.  Question:  But does this really prove anything, because of the effect of buoyancy?
• It is easy to compress a gas:  We took a syringe with plunger and stopper, removed  the plunger, and placed the stopper on the small end of the syringe.  We pushed the plunger back into the syringe, and compressed the air inside quite easily.
• Gases fill their containers completely:  We simply blew up a balloon to illustrate this point.
• Different gases can move through each other rapidly (diffusion):  Sarah described spraying perfume through an atomizer in one part of the room, and the students raising their hands when they smelled  it.  Those who were closer to the spray would smell it first, etc.
• Gases exert pressure.  We squeezed an inflated balloon, and could feel the pressure exerted by it.
• The pressure of a gas depends on its temperature: (thought-experiment suggested by Pat Riley)  Submerge the balloon in hot water, and then put it into an ice bath.  You could actually see the balloon expanding and contracting.
Very interesting, Sarah!

19 February 2002: Karlene Joseph (Lane Tech HS) Gas Laws
Karlene
handed out two flasks (round bottom Florence flasks, 500 ml).  Each flask contained 50-100 ml of water, with a partially inflated balloon inside.  Karlene asked us to explain how she got these balloons inside the flasks, and why they would remain there.  There were suggestions that the procedure for implantation of the balloons inside the flask might involve some sort of heating.  At this point she reviewed the gas laws:

Properties of an Ideal Gas
 Charles Law Volume / Temperature  = Constant Fixed Pressure Boyle's Law Pressure ´ Volume = Constant Fixed Temperature Gay-Lussac's Law Pressure / Temperature = Constant Fixed Volume Ideal Gas Law Pressure ´ Volume / Temperature = Constant General

Karlene then set up the apparatus, and we learned how to get the balloon inside the flask. She began by partially filling the flask with water, and heating it to boiling on a hot plate.  The air inside the flask was largely replaced by steam in this process.  After removing the flask from the hot plate, she quickly inserted an unfilled balloon into the flask, attaching its open end around the "lip" of the flask, so the un-inflated balloon lay along and inside the top neck of the flask.  As the flask cooled and steam condensed, the gas pressure inside the flask became less than air pressure outside, and the balloon began to inflate inside the flask.   The whole process took a few minutes.

At the suggestion of visitor Fred Schaal, we then heated the flask (with the balloon already inflated inside it), and again we produced steam from water.  The net effect was that the balloon became everted,  and it even inflated a bit outside the flask.  When we cooled the flask again, the balloon was pulled back inside the flask and filled with air. Note:  Here is the description of a very similar experiment:

Balloon in a Bottle:

Source:  Cool Science Demos from the Institute of Chemical Engineering [ICE] Workshops, University of Northern Colorado:  http://fileresource.sitepro.com/Applications/com/sitepro/modules/FileUtility/coolscience/1.pdf

Karlene then related these balloon phenomena to the Ideal Gas Law, P V = n R T.  When the balloon is put over the flask with trapped steam inside, the pressure P inside is the same as that of the outside air.  As the gas inside cools, steam condenses [number of moles n decreases], and pressure on the inside is reduced with decreasing temperature T.  The greater air pressure outside (in the room) pushes the balloon into the flask and inflates it. When the flask is heated again, steam is again created [number of moles of gas increases]; thus the pressure increases, pushing the balloon outside the flask.

Another Five (*****) Star job by Karlene!

05 November 2002: Tyrethis Penrice [Oak Park Elementary Schools]        FULL OF HOT AIR
Tyrethis told us we were full of hot air. Actually, that was the title of her presentation (though she may have been thinking that about us!).  She asked us some questions about balloons and hot air to get us started, accepting any answer because she was just warming us up for the main act! Tye placed a balloon over the neck of a 2 L pop bottle,  then poured hot water over a small part of the bottle. We could see the balloon inflate slightly, and then deflate when she used cold water in the same way. Gary cranked things up a bit by half immersing the bottle in the hot water in the coffee urn, and he got the balloon up to about 4 inches in diameter (again reversing when it cooled down). It was explained that the expansion of a gas when heated is due to the fact that the molecules move faster, which not only causes more collisions per second on the walls of the container, but also gives more push per collision.  In a related activity, several of us did a  "hot hand Luke demonstration, in which the screw cap on a small pop bottle was moistened and placed upside down over the neck opening.  We then placed our hands gently around the bottle (no squeezing, over there!), and we were rewarded by the cap doing a little dance as the expanding air forced its way through the thin film of water to escape. Thanks, Tye, for showing us how much can be taught with simple, inexpensive everyday items,

09 September 2003: Ben Stark [IIT, biology]       Calculating the Oxygen Content of Air
Ben
showed us a simple method to calculate the amount of oxygen in air, which also demonstrates the need for oxygen in air to support combustion.  He placed a candle upright into a shallow dish containing a little water, and put an inverted beaker over the candle, in such a way that the mouth of the beaker was completely underwater.  He marked the initial water level of the beaker, and determined V1, the volume of air in the beaker, as the total volume of the beaker, VT, minus the volume of water initially in the beaker, minus VC, the volume of the candle above the water level.  He obtained V1 = 310 ml. He then removed the beaker, lit the candle, and replaced the beaker.  As the candle burned, the level of water inside the beaker gradually rose.  After the candle flame went out, he again measured the volume of air in the beaker, obtaining V2 = 287 ml.   He then calculated the ratio V2 / V1 = 0.93.  Ben next used the perfect gas law, P1V1 = n1 RT1 and  P2V2 = n2 RT2   along with the fact that the pressure and temperature should be about the same before and after:  P2= P1  and T2= T1.  Thus, n2 / n1 = V2 / V1 = 0.93.  Therefore, in the process of consumption there has been a 7% loss in the number of moles of gas.  How come?

In burning wax, a hydrocarbon with a string of CH2 monomer units, the basic (approximate) chemical reaction is

2 C H2 (wax) + 3 02 (gas) ® 2 C02 (gas) + 2 H20 (liquid)
In other words, we convert 3 molecules of oxygen gas into 2 molecules of carbon dioxide gas.  So that the reduction in the number of oxygen molecules is three times the net reduction in the number of gas molecules. Thus, we estimate that 21 % of the molecules initially in the air were oxygen molecules expended in the process of combustion. This result is amazingly accurate!

Note: One must measure the volumes of the beaker, water, and candle carefully both before and after the candle burns to get precise results.

A breath of fresh air for us all!  Thanks, Ben!

09 March 2004: Ben Stark  [Illinois Institute of Technology, Biology]         Calculating the Oxygen Content of Air
Ben
repeated his lesson given at the HS Biology-Chemistry SMILE meeting of 09 September 2003.

Even better, Ben. Thanks!

09 March 2004: Bradley Wright [Eisenhower HS Blue Island, Chemistry]         How Do You See a Gas?
took us on an educational trip showing how to "see" a gas.  He brought out two "beakers" constructed from clear plastic pop bottles by cutting off the top portions.  In the first beaker, he mixed baking soda and vinegar, to produce bubbles of C02 gas.  Kids usually say "so what?", but Brad lit a candle, and lowered it into that beaker.  The flame went out!.  Then he removed the candle, lit it again, and put it  into the second beaker (filled with air). The candle flame continued to burn. Brad removed the candle from the second beaker. Next Brad poured C02 gas from the first beaker into the second beaker. Then he demonstrated that the candle would burn in the first beaker, but not in the second one.  Brad then produced a fresh batch of C02 in the first beaker.  He then poured the C02 down a V-shaped inclined ramp about 50 cm long, placing the lit candle at the bottom of the angle bar.  The flame in the candle went out -- again!  Wow!

Brad then demonstrated the Schlerien effect  --- which leads to wavy views when one looks down the highway or across the landscape on hot days.  Brad produced C02 in beaker 1 as before, and poured it into beaker 2 as before --- with the overhead projector lighting both beakers from behind.  The projected image on the screen became wavy as pouring took place.  We could literally "see" the CO2 gas being poured from one beaker to the other, by looking at the screen!

For additional details see the Bill Beaty's Weird Science [http://amasci.com/weird.html] article Threadlike Streams of Electric Windhttp://www.amasci.com/weird/unusual/airthred.html, or  and http://edgerton-digital-collections.org/techniques/schlieren.

12 April 2005: Walter Kondratko [Steinmetz HS]          Boyle's Law
Walter had a large syringe (with a plunger that could be used to adjust the volume of air in the syringe) attached to a pressure gauge by a tube (and the entire system sealed) and mounted on a clear plastic mounting so that we could lay it flat on an overhead projector, and project it on the screen. Walter then reminded us of Boyle's law PV = nRT (Pressure ´ Volume = number of moles of gas ´ gas constant ´ Temperature), so that at constant temperature for a given number of moles of gas, P ´ V = K, a constant (at least for modest pressures, where, for example, attractive forces between closely packed molecules do not come into play). This also means that P =K / V.

As we adjusted the plunger we obtained the data given below:

 Volume V: (mL) Pressure P: (PSI) P ´ V 24 15 360 19 19 361 14 25 370 11.5 31 357 29 12 348 34 10.5 357

Plotting the data as Pressure versus Volume, we obtained a (concave) curve that corresponded to P = K / VThe Gas Law works! Thanks, Walter!

24 January 2006: Ben Stark (Professor of Biology, IIT)           Measurement of the oxygen content of air
Ben
repeated a demo done at the Biology Chemistry SMILE meeting of 09 September 2003.  It is a version of the classic experiment in which a candle is burned in a saucer containing water with a beaker over it, and the water rises inside the beaker. Here is a copy of that description:

Ben showed us a simple method to calculate the amount of oxygen in air, which also demonstrates the need for oxygen in air to support combustion.  He placed a candle upright into a shallow dish containing a little water, and put an inverted beaker over the candle, in such a way that the mouth of the beaker was completely underwater.  He marked the initial water level of the beaker, and determined V1, the volume of air in the beaker, as the total volume of the beaker, VT, minus the volume of water initially in the beaker, minus VC, the volume of the candle above the water level.  He obtained V1 = 310 ml. He then removed the beaker, lit the candle, and replaced the beaker.  As the candle burned, the level of water inside the beaker gradually rose.  After the candle flame went out, he again measured the volume of air in the beaker, obtaining V2 = 287 ml.   He then calculated the ratio V2 / V1 = 0.93.  Ben next used the perfect gas law, P1V1 = n1 RT1 and  P2V2 = n2 RT2   along with the fact that the pressure and temperature should be about the same before and after:  P2= P1  and T2= T1.  Thus, n2 / n1 = V2 / V1 = 0.93.  Therefore, in the process of consumption there has been a 7% loss in the number of moles of gas.  How come?

In burning wax, a hydrocarbon with a string of CH2 monomer units, the basic (approximate) chemical reaction is

2 C H2 (wax) + 3 02 (gas) ® 2 C02 (gas) + 2 H20 (liquid)
In other words, we convert 3 molecules of oxygen gas into 2 molecules of carbon dioxide gas.  So that the reduction in the number of oxygen molecules is three times the net reduction in the number of gas molecules. Thus, we estimate that 21 % of the molecules initially in the air were oxygen molecules expended in the process of combustion. This result is amazingly accurate!

Note: One must measure the volumes of the beaker, water, and candle carefully both before and after the candle burns to get precise results.

A breath of fresh air for us all!  Thanks, Ben!

The result is remarkably accurate -- particularly in the light of the following effects:
1. Incomplete combustion
2. Production of Carbon Monoxide: CO
3. Dissolving of CO2 in H2O