You’re Getting Warmer! [W/Video]

The Little Shop of Physics has developed a series of videos called Flash Science, which show some exciting experiments that can be done with everyday items to demonstrate physics principles in a unique way. All of these experiments have been designed to be done by trained adults using proper safety equipment.


In physics, heat is something you do; it’s a verb. It is defined as the thermal (non-mechanical) transfer of energy. When you heat an object, you transfer energy to it, which can raise its temperature or even cause a phase change. Traditionally, three sources of heat transfer are cited: convection, conduction and radiation.


Radiation is the transfer of thermal energy using electromangetic waves, which includes visible light, infrared radiation, ultraviolet radiation, x-rays and microwaves and radio waves. A camera flash is designed to give off a whole lot of visible light in a short amount of time. The black ink in the newspaper absorbs this radiation and increases in temperature, while the blank paper reflects the light and does not warm up nearly as much.

Conduction and Convection

When a flame is held underneath a balloon, it’s no surprise that the balloon pops. The flame is at a high enough temperature to heat and melt (or even burn) the balloon, and the air under pressure inside quickly escapes. However, when the balloon is filled with water, the flame no longer pops it. The balloon is very thin, and the thermal energy quickly gets conducted to the water on the inside. The water has a very high heat capacity, so it takes a large amount of energy to increase the temperature of the water.

The water is also effective transferring the thermal energy away from the flame. The water will undergo convection; the warm water by the flame will move upwards, and be replaced by colder water coming in from the sides. Also, since water evaporates at 100°C, liquid water has a limit on how high of a temperature it can reach.


Evaporation is an extremely important and sometimes overlooked form of thermal energy transfer. Evaporative cooling is the mechanism behind human sweating, and the energy stored in evaporated water is extremely important in the Earth’s weather system.

In this video the flame hounds are soaked in a mixture of rubbing alcohol and water. While the alcohol burns, and releases thermal energy, the water evaporates and takes much of that thermal energy away from the flame hounds, so that it does not burn!

If you’re careful, you can even hold flaming bubbles in your hands!


Running electricity through the graphite pencil-lead causes the tip to get extremely hot, so hot that the graphite vaporizes and the vapor ionizes. These hot ions are used to cut aluminum foil, similar to how a plasma cutter or arc cutter works.

Erasing With Heat

Some erasable pens use thermochromic ink, which changes colors from dark to light when it is heated. When the ink is cooled (such as through the evaporation of a liquid), the ink becomes dark again. With this ink, you can erase and re-write messages over and over again!


Newspaper Radiation! [W/Video]


Radiation is the transfer of thermal energy using electromangetic waves, which includes visible light, infrared radiation, ultraviolet radiation, x-rays and microwaves and radio waves. A camera flash is designed to give off a whole lot of visible light in a short amount of time. The black ink in the newspaper absorbs this radiation and increases in temperature, while the blank paper reflects the light and does not warm up nearly as much.

Let us know in the comments below if you think this demo is something you would use in your classroom.

This video was produced by The Little Shop of Physics at Colorado State University in partnership with GE.


DIY: Erasing With Heat [W/Video]

The thermochromic ink used in erasable pens has a unique quality; it changes color with temperature. The molecules in the ink are oriented in layers, and when light passes through, the wavelength with the greatest constructive interference is reflected back. When applying heat in excess of 140℉ to the ink, the color changes form dark (black) to light (clear – appears to disappear). The change in temperature changes the spacing between the layers of molecules in the ink which changes the reflected wavelength leading to different colors. When the ink is cooled (such as through the evaporation of a liquid), the ink becomes dark again. With this ink, you can erase and re-write messages over and over again! NOTE: Please pay attention to the safety concerns expressed in the video when using as a demo.

Let us know in the comments below if you think this demo is something you would use in your classroom.

This video was produced by The Little Shop of Physics at Colorado State University in partnership with GE.


DIY: Baby Plasma Cutter [W/Video]

Pencil lead and some batteries make a small plasma cutter that is used to etch a pattern in aluminum foil.
In this simple but cool demo, you are able to observe how a high velocity ionized gas (plasma) conducts electricity across a small gap between the tip of the pencil lead (graphite) and a conductive solid (aluminum foil). When electric current is applied, the carbon atoms in the graphite vaporize and ionize creating a small ball of plasma that heats and melts the aluminum foil.

Let us know if you think this demo is something you would use in your classroom. See comments below!!!

This video was produced by The Little Shop of Physics at Colorado State University in partnership with GE.


Behind the Scenes with Light & Color: 10 Great Demos


Light and Spectrum is a common topic among all of the sciences. You will find a chapter devoted to it in Astronomy, Chemistry, Physics, and Biology. Therefore, enhancing your ability to teach this topic is going to benefit every member of the science department.


The RSpec-Explorer empowers teachers to have their entire class to experience quantitative spectroscopy at the same time and in a meaningful way. Up until now, it was very difficult to manage to get more than one person to be sure they were seeing the same thing through a diffraction grating or a refraction table. But with the RSpec-Explorer you can easily point out features in a gas tube line spectrum, a sodium lamp, or anything you can think of.

a lemon, an apple, and a green pepper are being studied using the RSpec-Explorer

In this picture, a lemon, an apple, and a green pepper are being studied using the RSpec-Explorer. The yellow lines in the left picture indicated that the apple is currently being investigated. The graph on the left shows that the apple is lined up (on the yellow line) and that its spectrum is mostly red, with a peak around 625nm.

In this article, I provide 10 examples of experiments you can do on light and spectrum, all of which are made easier by using the RSpec-Explorer.

1. Experiments on Color

white light of the iPhone flashlight turns out to be deficient in the light blues

The white light of the iPhone flashlight turns out to be deficient in the light blues. Unlike the even spread of color that would come from sunlight.

One of the first experiments you should do is to demonstrate that white light is made of colors. The term “white” is often used by scientists to refer to a light source that emits or reflects all visible wavelengths (400-700nm). However, the human eye cannot distinguish this real white light from a light source that is made of only a few colors. For example, if you examine a cell phone flashlight feature through a diffraction grating (such as the one on the RSpec Explorer’s camera) it will reveal that this apparently “white” light is actually missing some of the deep blues. Also, if you look at a “white” fluorescent lamp tube, it will reveal that it is made of several distinct colors but not a broad spectrum (like say for example a sunbeam, or a white incandescent lamp).

If you have a color mixing device (three colored lamps would work, or three lamps with filters) you can demonstrate to yourself that “white” light can be created by mixing Red, Green, and Blue (RGB). This is how a cell phone screen makes white light, and a computer screen, and most projectors! There are many sources available for this experiment.


Magenta being faked by mixing red and blue. The pinkish looking light is diffracted revealing it is a mixture of red and blue. The spectrum on the right reveals the peaks of these two colors: 425 and 610nm.

A better trick is to mix just two colors and get a new color that will completely fool the eye. A major example is to mix red and green light and make “yellow.” I put the quotes here again because it only appears yellow – there is nothing yellow about green and red light mixed, except that it can fool the eye by appearing to be yellow. Mixing red and blue light makes “magenta” light and mixing blue and green makes “cyan” (again the quotes describe the appearance not the reality of the light).

The advantage of analysis through a diffraction grating is that it can easily discern the two colors, which diffract differently. The spread or “dispersion” of the light is linearly-dependent on its wavelength (to a good approximation). That’s how we can separate the light by its wavelength and reveal whether we are looking at a true color or only a synthesized one.

2. Ionized Gases

Ionizing Gases to display their spectra is an important activity in most science classes. Of course, you want to point out that different gases have different spectra and these can be used for identification. Every noble gas was identified first based on its spectrum. (How else could you tell Neon from Argon, seeing as how they are both chemically inert?!!)

Gas tubes discharging

Gas tubes discharging: Hydrogen, Helium, Neon, Mercury.

The most important example is hydrogen, which is not a noble gas, but which has a readily recognizable spectrum. The Balmer Series (n=2) is the visible portion of the spectrum. It has a very obvious and bright cyan (486nm) colored line, a somewhat less bright red line (656nm), and a few violet lines (434nm, 410nm). Invisible is the Lyman (ultraviolet, n=1) and the Paschen (infrared, n=3) series. The hydrogen spectrum is important, not just because it is so familiar, but because it can be calculated easily (using the Rydberg formula), and it was also the spectrum that was used by Niels Bohr when he applied quantum theory to explain atomic spectra for the first time.

The Balmer series

The Balmer series for hydrogen contains the four visible lines of hydrogen’s spectrum and all of these transitions involve the n=2 orbital (marked in yellow).

The Rydberg Formula

The Rydberg Formula for Hydrogen’s spectral wavelengths

Helium, on the other hand, is not as familiar but can be made so by learning to recognize it by its bright yellow (589nm) line. Also, the story that the helium absorption lines were first seen in a solar eclipse is a good history of science tidbit. That is how helium got its name, from the sun god – Helios. Also, helium looks yellow-pink when ionized, where hydrogen usually looks red-purple.

four characteristic spectral lines

The four characteristic spectral lines in the Balmer series for Hydrogen.

Recognizing Helium

Learning to recognize Helium based on its very bright yellow discharge line and its pale yellow spectral tube.

Neon is amazing to look at even without a diffraction grating. Its pastel-electric red glow earned it its name as the “new” element for the electric age. The diffraction grating reveals that it is saturated with reds, yellows, and a few scattered greens.

In the plasma globe you can find ionized gases and with the RSpec-Explorer and (if you can line it up carefully) you can identify the gases inside (helium is the main one).

paper clip bent

A paper clip bent to include a small loop does a great job of carrying salt to the flame.

It is also possible to burn salts and reveal the spectral lines.  The most obvious salt is table salt, sodium chloride, which burns well in a paper clip loop held over a candle flame. Teachers often dissolve a lot of salt in a little water which can sometimes help (dissolved ions have more surface area/volume than crystals so they burn easier). The yellow sodium “doublet” (two very close wavelengths at once) easily identifies it. You can also recognize sodium in yellow streetlights at night. Other salts that emit good colors will contain copper, strontium, calcium, potassium, and iron. Which you can usually find in the chemistry storeroom. All of these are often used in fireworks (usually mixed with magnesium and gunpowder) and if you need help getting the fire hot enough, you should try dissolving them in methanol. Safety first! Be sure to have safe water on deck for emergencies and a fire extinguisher is a good idea, too.

3. Investigate Different Light Bulbs

These days people are very interested in how all the different types of light bulbs make light. Diffraction is the best way to identify how the light is made. If you look at an ordinary incandescent bulb you will see it has a broad spectrum with a lot of yellows and reds giving it a “warm” glow. On the other hand, fluorescence light bulbs contain mercury and will have several easily recognizable spectral lines that correspond to that element. Mercury is a good choice for fluorescence because the many energetic purples and UVs in the spectrum can give energy to fluorescent paints which reradiate that energy as visible light. If anyone doubts that there is mercury in our light sources, they should be easily convinced by this demonstration!

mercury discharge tube

A mercury discharge tube demonstrates several violets and greens. But very little red and yellow. Invisible is the UV.


fluorescent light bulb

A fluorescent light bulb shows many of the same spectral lines revealing that it contains mercury.

Compared to incandescent bulbs, fluorescent bulbs tend to make people look drained of color. This is because the high amount of blues and purples can cast an “unhealthy” purple glow on you. A 100W incandescent bulb nearly imitates the sun’s spectrum. Which peaks in yellow-green, giving you that healthy glow.

Incandescent Bulb

An incandescent bulb reveals its warm colors by peaking in the reds and yellows.

You can also investigate other light sources such as white diodes (which have a lot of purples because they fluoresce, too), yellow sodium parking lot lights, or even a plant light. Plant lights aim to provide the two spectral colors of photosynthesis – blue and red. Green plants reflect green light and thus they do not absorb it for making glucose. Red light also provides a signal to the plant to let it know that the day is long enough (i.e. spring or summer) to start investing itself in growth. (Some plants actually suppress their growth in summer to take advantage of a less competitive winter season.) Anyways, plant lights provide these non-green colors in high supply.

4. Analyze the Wavelengths of Lasers and Diodes

Light Emitting Diodes are a ubiquitous source of light in our lives. In most cases, diodes will be sold to emit a specific wavelength of light but in actuality, there will be a spread of color about this “nominal” value. (Nominal is an engineering term meaning “named” or expected, as opposed to what actually results during the experiment.) For example, a “626nm” LED might emit 96% of its light between 610 and 632nm. This amount of spread can be measured by the RSpec-Explorer and it’s interesting to compare this with laser light.

orange diode

The orange diode demonstrates that its wavelength is quite spread out over the 30nm that surrounds its nominal value.

A HeNe laser

A HeNe laser demonstrates both that it is monochromatic and that it contains neon by emitting the 626nm red that helium lacks.

Lasers have “monochromatic” light. This means that it is very nearly only one specific wavelength. These wavelengths are usually listed on the laser itself. A good experiment would be to verify that the wavelength printed on the laser is actually the wavelength it emits.   Even diode lasers are usually quite monochromatic. “Lasing” requires the light to be nearly one wavelength – lasers are a good example of light standing waves.

helium neon laser

A bare helium-neon laser glows yellow pink with helium but emits a neon red beam.

When it comes to red lasers there are many different types. Helium-Neon lasers will have different wavelengths than red diode lasers. You can use the RSpec-Explorer to prove that it is actually neon that emits the red light in the helium-neon laser. This is a good demonstration of the power of spectral analysis to identify elements. If you have a bare helium-neon laser it can be particularly engaging in this activity.

5. Investigate Fluorescence

tonic water

A violet laser energizes the quinine in tonic water.

Fluorescence is always an engaging activity.   Energetic light (such as UV or violet) lands on a substance that can absorb it and that energy is re-emitted as less energetic visible light. For example, a black (UV) light might shine on your socks (which have fluorescent detergent) in them and then white visible light will be emitted.

Good candidates for investigation with the RSpec-Explorer include tonic water, highlighters, extra virgin olive oil, Willemite, and phosphorescent vinyl sheets. All of these will glow under UV or violet light (such as a violet laser or black light) but the olive oil works better with a green laser (the yellow olive oil absorbs violet light very quickly).

light off

Above: UV Light is turned off. 

Lights On

A few fluorescent rocks with UV turned on. Willemite is in the middle.

Phosphorescence is a special type of fluorescence in which the emission of the light is suppressed for an extended period of time (the atomic transition is slower). In fluorescence, the emission of light is nearly instantaneous. In either case, the wavelength of the emitted light is always longer – less energetic. The words phosphorescence and fluorescence are only historical. Not all phosphorescent materials contain phosphates (though most do) and not all fluorescent materials contain fluorides (though many do, including toothpaste). Willemite, which is a fantastic glow rock, contains neither phosphorus nor fluorine.

6. Measure Temperature Using the Blackbody Curve

Turn on your electric oven, toaster, or electric stove and it will first glow red-hot, then yellow-hot, and if we went further it would glow white-hot. This change in color with temperature was described mathematically by Lord Rayleigh, James Jeans, Wilhelm Wein, and finally Max Planck. The Rayleigh-Jeans Law described the long wavelengths and Wein’s Law described the short wavelengths, but both “laws” failed outside of those conditions. Planck was the first to solve the emission problem for ends of the curve. Planck’s function is also called the “blackbody” curve because even a black object will be seen to emit light in these proportions if it gets hot enough.

We can use this curve to determine how hot our light bulbs are. The problem is that most of the light they emit is infrared which is generally not visible. If you are willing to accept that there is an enormous amount of invisible light, then you will be able to approximate the temperature of a glowing hot object by this method. You may be surprised to find out that even little circuit-lab style bulbs are actually heating up to about 4000K – but this is consistent with theory.

Reference Library

The Reference Library includes Planck Curves which can be used to fit to our spectrum. Hotter objects have steeper slopes on the left side. The right side is incomplete due to invisible infrared light.

I am not saying that you can get a highly accurate measurement with the RSpec-Explorer camera, but you can approximate the temperature reading, and probably within 15% (be sure to turn the brightness down in the settings). It is impressive that lightbulbs get this hot to glow. It also helps us appreciate why they must be contained in bulbs – if they were exposed to the oxygen in the air at these temperatures they would immediately burn up and break the circuit.

It’s fun to compare these light bulb filaments to the temperature of the sun which is a G2V star (there is a star reference collection in the References, too). The temperature of the sun can be determined from the black body curve as well. In fact, this is how we measure the temperature of the sun – at least on its surface!

Measuring temperature using the blackbody curve is a good way to get Modern Physics concepts into your classroom.

7. Diffraction Experiments

In Astronomy and Physics, the idea of Diffraction is a commonly taught subject. Diffraction of light is one means by which we can separate it based on its wavelength. A diffraction grating is made when a laser cuts tiny grooves into the surface of a piece of plastic or glass. A good example of one is a CD.

Some themes of diffraction are that the smaller the distance d between the grooves, the more dispersion, X, you will get. The light will spread apart further from its straight line path. Also, the longer the wavelength, λ, of the light the more easily you can disperse it. And of course, the more space it has to travel before it lands on a screen the more it will disperse. This length is usually called L (the distance to the screen or camera). All of these ideas come together in the diffraction formula:


Xm = m λ L / d


where m is an integer, usually 1, that tells you the “order number.” We need m because the pattern will repeat itself about twice as far out, and that is called the 2nd order. Usually, the 2nd order is much less bright than 1st order diffraction. This formula is an approximation, assuming that the light is not being diffracted at large angles from the straight-ahead path, it works well for angles under 30 degrees (ie first order).

measure both

You’ll have to measure both the dispersion X (left) and the distance to the camera L (right) if you wish to apply the diffraction formula. The wavelength λ is given in this case, which is unusually convenient. What is not given is the groove spacing d.

A good lab would be to use this formula to try to measure the line spacing “d” for the diffraction grating of the RSpec-Explorer camera. The units should be in meters/groove or meters/line. Use meters as the unit for X, λ, and L.


8. Measure the Wavelength of Infrared

The wavelength of Invisible Infrared Light can be measured with a diffraction grating and a digital camera (which can see the infrared light). But, the RSpec-Explorer makes this easier because it can tell you the wavelength based on the distance the light is diffracted on the video screen. A TV remote control can provide a source of near-infrared light (should be between 800 and 1000nm) but I have had more success with loose infrared diodes because I can crank up the voltage and get them to shine very bright. To ensure success, have a very dark room with the diode close to the camera. Also, when you rotate the camera you should be able to see the first order diffraction of the infrared light.

infrared diode

An infrared diode (bright circle) and its lens flare (dimmer circles) are plainly visible in this screen capture. Lens flares occur when bright lights are improperly focused by lenses. Like all cameras, the RSpec-Explorer demonstrates this feature.

To perform this experiment, get your diode showing up very bright, and line it up on the yellow calibration line. You will not be able to see the diffracted light because it will be diffracted so far that it will not appear on the screen. So instead we will have to recalibrate the camera to see further than the visible spectrum. First, line up a second light source with a familiar visible spectrum such as a Hydrogen gas tube or a white diode. Then, once you have it on the yellow line, rotate the camera to the right. This should cause the two light sources to be off-camera to the left but the color spectrum will remain. Based on your knowledge of that familiar spectrum you can recalibrate the camera. Go to Tools à Calibrate à Linear. Click on the first familiar pixel and enter its wavelength (in the case of hydrogen this would be the 486nm cyan line). Then click on the second pixel and set that to another wavelength (in the case of hydrogen this would be the 656nm red line). Now “apply” that calibration and close the window. Turn off your familiar light source and the infrared light should now be visible around 900nm. Be sure to move the yellow bars to sample it and explore its peak brightness. Like all diodes, infrared ones have a significant spread.

explore the infrared

A screen capture took after recalibrating the camera to explore the infrared. In this case, the hydrogen spectrum was used as a familiar reference. The next step would be to double check that the infrared diode was lined up with the hydrogen tube at the origin, and then turn off the hydrogen tube to reduce accidental light contamination.


9. Experiments that use the Intensity Feature 

Crossing Polarizers

Crossing Polarizers can reduce the Intensity of the light that comes through.

You probably have already noticed that brighter spectral lines show up as higher peaks on the y-axis. This is particularly obvious with the hydrogen spectrum’s cyan line at 486nm which is much brighter than its red 656nm line. This intensity reading feature can be helpful in other experiments as well. To take advantage of this feature, be sure to turn off the “Auto-Scale Y-Axis” feature on the bottom right panel of the screen.

A good experiment to try out is to block a light source with two polarizing sheets. When the polarizers are rotated they block more of the incoming light. As the relative angle increases (to 90 degrees) the blocked light source dims to nearly zero transmittance. This reduction in brightness is supposedly dependent on the square of the cosine of the relative angle between the polarizers. This intensity function I(θ) = Imax cos2(θ) is known as Malus’ Law. It is helpful if the source of light is monochromatic.

Another experiment that you can try (one that probably belongs in a Chemistry class) is that a higher concentration (molarity) of dye will block a proportionately larger amount of light. This is known as Beer’s Law. It might be best to start with a clear water sample for calibration then slowly add dye. I have had a lot of success with Coca-Cola. Again, it is helpful if the source of light is monochromatic.

Beer’s Law Experiment

In this Beer’s Law Experiment, the concentration of the solution is increased, causing the intensity of light to be decreased. Moussing over the central plateau makes an intensity measurement. It is important to set a constant scale for the y-axis. I have chosen not to fill the graph with color because the wavelength of light is not being measured, only brightness.


10. Astronomy Experiments

The RSpec software that is employed with the Explorer camera was originally produced to serve astronomers. Thus, there are many vestigial traces of this in the reference libraries and in the training videos that accompany the device. It can be fun to take advantage of these features and see how far one can push the camera.

reference library

Here the reference library is used to investigate a g2v star. This is the same type of star as the sun. The next move would be to click Planck and check that 5700K is the right temperature for this curve.

You can view the solar spectrum by reflecting sunlight from along the length of a needle at the camera. Since all that is required is to view it is a bright source of light lined up along the yellow 0 nm line, it is quite possible to take advantage of this and observe the sun. Most visible in the spectrum will be the g2v black body curve and, if you zoom in, the Fraunhofer Absorption lines.

Perhaps an easier demonstration is to view the clear, blue sky through a slit. This reveals that the blue sky is actually a mixture of all colors with more blue than any other color. The truth is that there is actually a little more violent than the camera can reveal but like the human eye the camera is less sensitive to violet than to blue. This “unsaturated blue” (meaning blue + white) is consistent with the Rayleigh Scattering Model for why the sky is blue.

If you own a telescope you may be able to analyze the spectrum of the stars with the RSpec-Explorer. Because of the sensitivity limitations of the camera, it is not possible to observe stars without a telescope. But, a telescope can help a lot. Be sure to know which star you are looking at (I recommend Sirius, Betelgeuse, and Vega), then look up what type of star they are (a1v, m2i, a0v respectively).



spectrum analyzing equipment

Old-fashioned – but not obsolete – spectrum analyzing equipment.

Experiments on light can be very engaging, but they can also be very confusing. It is important that we take steps to ensure that our students are able to view what they are supposed to be seeing, and recognizing what is being pointed out. The RSpec-Explorer projected overhead for your students’ benefit is probably the best way to in engaging your students in spectroscopy (especially if used in conjunction with hand-held spectrum analyzing devices). I have found that students are very interested in cameras and how they can see things that our eyes cannot. If building a community of learners in your science classroom is your goal then you should add this device to your collection of lab equipment.

Projecting Image

Projecting this image for your class will ensure that they will understand what you mean when you say, “Yellow light is dispersed further than violet light, and red light is dispersed the most.”


James Lincoln

Tarbut V’ Torah High School

Irvine, CA, USA

James Lincoln teaches Physics in Southern California and has won several science video contests and worked on various projects in the past few years.  James has consulted on TV’s “The Big Bang Theory” and WebTV’s “This vs. That” and the UCLA Physics Video Project.

Contact: [email protected]


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Recreate Physics History: Build a Voltaic Pile

In the late 1700s, Italian scientist Luigi Galvani stumbled across one of the most important discoveries of all time. He found that frog legs would contract when some of the muscles and nerves were connected – even when the frog was dead! Galvani attributed this phenomenon to the idea that animal and human brains produce electricity, which he referred to as “animal electricity.” He surmised that this electricity was then stored in the animal’s muscles after being transported to the muscles through the nerves, much like how electricity is stored in a Leyden jar. According to Galvani, when certain muscles and nerves were connected, animal electricity discharged and the muscles contracted.

Can’t view this in YouTube? Try watching in Vimeo.

Italian scientist, Allesandro Volta, read about Galvani’s work and through experimentation came to strongly disagree with Galvani’s explanation of the phenomena, especially the idea that living beings produced “animal electricity.” Volta instead believed that the electricity present in the frog was the result of the contact of the metal probes with the frog tissue. In 1791, Volta produced continuous electric current when he placed a cloth soaked in salt water between silver and zinc disks. In 1800, Volta discovered that the current increased when he stacked several pairs of these single electrochemical cells together. This device became known as the voltaic pile, and was the first electrochemical battery. You can learn more about the debate between Volta and Galvani and its significance to the fields of both electricity and anatomy by viewing the video below:

You can make your own voltaic pile out of simple and inexpensive materials. Although Volta used silver and zinc, it is more feasible – and inexpensive – to use copper and zinc for the metal disks. Even though pennies are no longer made of copper, their copper coating still makes them a great choice for copper disks, and zinc disks can be obtained by purchasing galvanized electrical boxes and punching out the holes. The electrical box seen on this page and in the video was purchased at a local home improvement store for only $0.74 and provides 17 zinc disks. Other materials needed include thick card stock, salt water, a voltmeter or multimeter, and scissors. Wooden dowel rods poked into modeling clay can provide vertical support for the pile as more and more cells are stacked on top of each other, and more closely replicates the design of the original Voltaic pile. You may also want to include an LED and connecting wires as your voltaic pile should generate enough power to light an LED.

Voltaic Pile Materials

Voltaic Pile Materials

To make the voltaic pile, cut out card stock disks the size of a penny and soak them in a cup of salt water. To make a single cell, place a card stock disk that was soaked in salt water on top of a zinc disk, and then place a penny on top of the card stock. Touch the positive probe of a voltmeter to the copper and the negative probe to the zinc and you will find that the electric potential difference, or voltage, of this simple electrochemical cell will likely be between 0.60 V and 0.80 V. If you make another cell and stack, or pile, it on top of the other so that you essentially have two cells in series, you should find that the resulting electric potential difference is between 1.20 V and 1.60 V. If you continue to pile single cells made of a zinc-card stock-copper sandwich on top of each other, you will find that the voltage increases with each additional cell. Small wooden dowel rods poked into modeling clay can be used to keep the stack of electrochemical cells from falling over. The voltaic pile illustrated below was made of multiple zinc-card stock-copper cells and had an electric potential difference of 3.26 V at the time of the photo, although it had initially read larger. Unfortunately, you can expect to have fluctuations and inconsistencies in voltage readings, but readings should generally increase as additional cells are added.

Voltaic Pile with Multiple Cells

Voltaic Pile with Multiple Cells

You should be able to connect an LED to your voltaic pile and watch it light. LEDs are directional, which means that the positive lead (the longer of the two leads exiting the bottom of the bulb) must be connected to the positive (copper) terminal of your voltaic pile.

LED Lit by Voltaic Pile

LED Lit by Voltaic Pile

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Exploring Matter: Chemistry Demonstrations

In this issue of CoolStuff, we’ll once again welcome guest author Patty Carlson. At New Trier High School, Patty is known for her ability to make chemistry come alive. Her flair for great demonstrations and labs certainly comes through in the upcoming collection of activities illustrating properties of matter. I know you’ll enjoy sharing these marvels with your students! I have known and worked with Patty Carlson for over a decade and feel privileged to teach and grow with her. Patty is an energetic, creative, and caring professional who has earned the respect of her students and the admiration of her colleagues. She relates well to students, knows her subject inside and out, and is enthusiastic about sharing her knowledge with others. Patty is intrigued by the simplest things, which, I believe, explains her success. She seems more fascinated with the natural world around her each day and she shares this growing sense of wonder with her students.

Exploring Matter Activities

Chemistry is all about studying matter and how it changes. Fortunately many characteristics of matter are macroscopic, that is, we can directly observe them without the aid of any lens other than those in our own eyes. We can watch as matter is mixed or reacted and ultimately may be able to infer something about its deeper, more abstract structure (electron configurations that determine bonding, or the pairing/ unpairing of electrons that result in the magnetic properties of certain elements, the molecular shape of a molecule that influences its polarity, or the “packing” of atoms that determines the density of a substance.)

In the following activities students are encouraged to poke, prod, pour and play (wow…..that’s a lot of alliteration…) with matter and watch as it responds. I’ve found over the years that using familiar, household items as much as possible reduces the intimidation factor that some students feel in physical science classes, especially in the beginning of the year when many of the concepts explored below are introduced. The Reddi-Whip will go fast as will any extra soda you might have leftover to wash it down. Make sure you have lots of Styrofoam cups too because even my high school students want to see that acetone/cup demo over and over and over and over…

Patty Carlson ~ New Trier High School, IL.

I. Like Dissolves Like

Ever have a nasty stain on your shirt that won’t come out in the wash, no matter how many times you try, and yet that same stained shirt comes back from the dry cleaner looking like new? If you have, you’ve experienced the chemical phenomenon of “like dissolves like”. That is, substances tend to dissolve in things that are similar to them. By ‘similar’ in this case we mean in terms of their polarity. Some stains dissolve better in a polar substance like water and some stains require a more non-polar substance to dissolve them away.

Let’s consider two solvents that are pretty different in their polarities in order to explore this topic. Water, which we said is a polar solvent, dissolves almost anything that is polar, such as salt and many other ionic compounds. Water can’t dissolve everything, though. Try removing fingernail polish with water and you’ll see what I mean. Acetone, a solvent with some non-polar properties, is commonly used to do that job. Acetone is an effective solvent for all sorts of non-polar substances.


Place two large glass beakers side-by-side. Pour water into the first beaker until it’s about half full. Place a Styrofoam cup in the water beaker. Nothing will happen. Styrofoam is non-polar, water is polar and, since “like dissolves like”, they will not dissolve in each other.

The goo you retrieve from the beaker is actually polystyrene plastic (#6 in recycling code) and is the same plastic used to make plastic table ware, etc. You can shape it any way you wish while it is wet and it will harden over time as all the acetone completely evaporates away. In order to completelydissolve the plastic, you’d need a stronger and more non-polar solvent.

Now pour some acetone into the other beaker and place another Styrofoam cup into that beaker. You’ll see the cup slowly break down until it is just a glob of goo. Acetone can get in between the components of the polymer of plastic and allow the air in the cup to escape (don’t worry, they don’t use CFC’s in Styrofoam anymore so there is no harm to the environment when doing this demo).
Place starch packing peanuts (the environmentally friendly packing option commonly used today) in a beaker of acetone. Since the starch packing peanuts are polar, they will not dissolve in acetone. Put the starch packing peanuts in a beaker of water, mix around a bit and you’ll see they dissolve readily

Starch peanuts in acetone

Old-fashioned Styrofoam packing peanuts are fun to play with too. You’ll need a large beaker filled about half full with acetone. Have someone ready with a large wooden spoon and start loading the Styrofoam packing peanuts into the beaker as your helper stirs like crazy. You’ll be amazed at how many peanuts will fit into the beaker.
II.  Density: 

An interesting and often surprising property of a substance is its density, or the ratio of a certain mass of that substance to its volume. As long as you keep the temperature the same, the density of a particular substance never changes. You may have felt how heavy a chunk of lead is compared to a chunk of aluminum of the same size or perhaps you’ve held a jar of metal mercury and been amazed by how heavy even a small amount of this element is. These are differences in density. Since chunks of lead and jars of mercury are a little hard to come by, let’s explore this with some pretty ordinary stuff: Coke and Diet Coke.

Get a large, glass beaker (or aquarium) filled with water, a can of Coke and a can of Diet Coke. Place both cans in the water. The Coke will sink; the Diet Coke will float. Ask students to hypothesize about why this is so. (Caption: The difference between the two densities is real, but subtle. Make sure to do this in a large volume container (1000-2000 ml) in order to make the difference as obvious as possible.  The density if Coke is slightly above 1.0 g/ ml and the density of Diet Coke is just about 1.0 g/ ml. The density of water (at room temp) is 1.0 g/ml. We assume the aluminum cans are identical in density.)
Activity/ Lab:
Challenge your students to design an experiment that will allow them to determine exactly what the densities of the two sodas are. This can be done easily using small graduated cylinders (10 ml) and an electronic balance. For example, they can pour 2 ml of coke into the graduated cylinder, place the cylinder on a balance and record the mass. (Of course, they should correct for the mass of the graduated cylinder.) This would be their first “data point”. They can repeat this technique with 4 ml, 6 ml and 8 ml of the Coke and corresponding masses for those volumes. This entire process is repeated with Diet Coke. Make sure students don’t get them mixed up. They may taste different, but they look identical in the lab.Once they have gathered their data, can find the density by one of two methods: graphing the data and finding the slope of the mass vs. volume line (most accurate), or simply finding the average density from the data points. When graphing, students should include 0,0 as a data point, since zero volume of soda has a mass of zero.The students will probably guess that the only real difference between these sodas is the sugar content. Coke contains approximately 39 grams of sugars (high fructose corn syrup and/or sucrose, which is regular old sugar) . Diet Coke contains Nutrasweet (aspartame) and since Nutrasweet is SO much sweeter than sugar, only about 100 milligrams per can are required to get it to match Coke’s level of sweetness. That’s a pretty big difference and the reason for the difference in densities of the two sodas.If you want to make it more interesting, try the new low-carb Coke, C2, and see where its density falls with respect to the other two. It contains a combination of artificial sweeteners (aspartame, acesulfame potassium, and sucralose, which is Splenda) in addition to high fructose corn syrup and/or sugar. You can also try different brands. Tab contains saccharin and Diet Rite uses a combination of artificial sweeteners, giving them a slightly different density.
III. Density Columns:
Here’s a Demo and/or Activity that uses the concept of “like dissolves like” and density! You’ll need: Dark Karo syrup, Water (with food coloring too help students identify which layer it is in the column), Vegetable oil, Rubbing alcohol (isopropyl alcohol), and Large glass cylinder (or any long tube will do. It doesn’t have to be graduated). To do this as a demo, take the glass cylinder and pour in the dark Karo syrup (the most dense in this list). Then carefully pour in the colored water. You’ll note that they mix a little bit (there will be a “blur” between the two layers), but they are still distinctly layered. (The sugary syrup has some polar properties and the water will dissolve it at the point of contact.) Then pour in the vegetable oil. Because oil and water don’t mix (oil is non-polar, water is polar) they will also form distinct layers. For the last layer, add the rubbing alcohol. This can get messy and the column will need time to settle itself down. The alcohol will dissolve in water (alcohol has a polar region), but the oil will form a barrier between the water and alcohol. When you pour the alcohol into the column, it will come into contact with the oil and go from clear to murky. Again, there will be a blurring of the “line” between the two layers due to partial solubility (rubbing alcohol has non-polar parts too and oil is non-polar so a little mixing will occur).
To do this as a lab activity, give students smaller columns and the same 4 liquids. Let them pour the liquids in any order they wish. Based on their observations, they should be able to figure out which liquids are more dense than which others. Finally, they will be able to rank the liquids according to their relative densities. Rubbing alcohol 0.87 g/ml Vegetable Oil 0.91 g/ml Water 1.00 g/ml Dark Karo Syrup 1.37 g/ml To add a little complexity to this activity, ask the students to infer the approximate densities of the following solids: Ball bearing, Plastic bead, Cork, Rubber stopperThey can do this by dropping the objects, one-by-one, into the column and see if they float or sink in a particular layer. If they know the numerical value for the densities of each of the 4 liquids, they can approximate a value for the density of each of the solids.  Students should observe the following sequence, in order from least to most dense: Cork – Rubbing Alcohol – Vegetable Oil – Plastic Bead – Water – Rubber Stopper – Karo Syrup – Ball Bearing
IV.  Classification of Matter:
Matter is anything that has mass and takes up space (has volume). We can separate the matter that we know about into two huge categories; mixtures and pure substances. Well, what are mixtures? Mixtures are physical combinations of at least two pure substances. Most of us are much more familiar with mixtures than pure substances and they are indeed much more common in our everyday experiences. For more on mixtures, check this out:
Mixtures can be further categorized into homogeneous and heterogeneous mixtures. Homogeneous mixtures are mixtures with the same composition throughout. Let’s say you stir some powdered Kool-aid mix into a pitcher of water. Once the powder is dissolved, doesn’t that Kool-aid look and taste the same from the first sip to the last?  Compare that to some orange juice with pulp in it. Let’s say your brother never shakes up the carton when he pours himself a glass of juice. By the time you get it, there is a huge blob of pulp at the bottom of the carton. Now your glass is a combination of juice and big globs of pulp. That, my friend, is a heterogeneous mixture and a gross one at that. Heterogeneous mixtures are not uniform in composition at all. Now, what are pure substances? These are either individual elements right from the Periodic Table or compounds (chemical combinations of those elements). The element Iron, for example, is a pure substance. Let’s say we let that iron sit around outside for while and we notice it starts to rust. It has undergone a chemical reaction and combined with oxygen in the atmosphere to create iron oxide, which is a compound, and, by itself, also a pure substance. These elements aren’t “mixed’ together like the mixtures we talked about before, they are BONDED together in a chemical way that won’t allow you to un-bond them very easily. Getting confused? Let’s take a look at some examples and maybe things will clear up…

Student Activity:
This is a station-based activity (or “smorgasbord” as my buddy Chris Chiaverina calls them) so you’ll need lots of lab space and a place for kids to walk around in small groups. Collect items like the following (and/or add your own!) and place them around the lab benches. Ask the students to:

1) Identify the category of matter:
a. Is it a pure substance? If so then is it an element or is it a compound?
b. Is it a mixture? If so, then is it a heterogeneous mixture or homogeneous mixture?
c. (Optional) Have the students write down the criteria they use for their categorization schemes.

Devise a separation strategy for any mixtures found. In other words, if you think you’ve spotted a mixture, how would you separate it into different components, and (if possible) all the way to the pure substances that comprise the mixture? (Remember, pure substances cannot be separated by physical means. They must be separated chemically, or, in the case of elements, by splitting atoms! That’s beyond the scope of the activity for the day.)

Suggested Items:
1. Aluminum foil: (Pure substance, element)
2. Lucky Charms: (Heterogeneous mixture.)  Separate physically. Visually identify the cereal from the sugary charms and manually sort into two piles. Separation beyond this is too difficult.
3. Orange juice with pulp: (Heterogeneous mixture.) Separate by gravity filtration of the pulp.
4. Salt water: (Homogeneous mixture.) Separate by boiling away or evaporating the water and leaving the salt crystals behind.
5. Salt, sand and water: (Heterogeneous mixture.) Separate by filtering out the sand, boiling off water and leaving salt crystals behind.6. Reddi Whip dessert topping: (Homogeneous mixture…really a colloid, but that may be too fine a point here.)  Separate gas from solid portion by heating it. Gas will bubble out since less soluble at higher temps, leaving solid portion behind. Separation beyond this is too difficult.

7. Oil and vinegar salad dressing:(Heterogeneous mixture.) Separate by difference in density.  (A separatory funnel is a good tool for this.)

8. Chocolate Silk Jif: (Homogeneous mixture.)Separation strategy: good luck J Some homogeneous mixtures are so uniform, even at the microscopic level, they seem extremely difficult to separate by conventional means.

9. Aspirin (make sure this is pure aspirin with nothing added, like buffers or anything else): (Pure Substance, compound – acetylsalicylic acid.  All aspirin is this compound. People buy different aspirin products for different reasons. Buffered aspirin helps those prone to stomach upset, etc…)
10. Juice Bar Candy Refreshee Spray Perfume:
(Homogeneous mixture.) Separate by differences in boiling point using fractional distillation. (What’s that? Check this out:http://chemistry.about.com11.Iron and Sulfur
(literally iron filings and powdered sulfur): (Heterogeneous mixture) Separate by difference in magnetic properties.

12. The Ink in a Black Sharpee Pen:
(homogeneous mixture.) Separate by chromatographic means. That is, the different pigments in the pen that make “black” can be separated based on their solubility in different solvents of different polarities.
Demo Idea:
Combine the iron and sulfur mixture from # 11 in a test tube.  Heat over a bunsen burner under the hood, you can show the students that a new substance is chemically formed. It is a pure substance and a compound, iron sulfide (FeS). It no longer has any magnetic properties at all and proves that is has been chemically changed from two elements to a single compound with completely different physical properties.

What’s the Matter?

A look at some weird solidy-liquidy type stuff:

We are all probably aware of the basic states of matter: solids, liquids and gases. When we see a solid, we expect it to act like a solid, that is, have a definite volume and a distinct shape (at a given temperature). When we see liquids we expect them to behave like liquids. They should flow easily, no matter how hard or gently we stir them around.

Are there substances that don’t behave the way we think they should? Sure! They’re called non-Newtonian substances.

Slime Receipt Things you’ll need:
Electronic balance
150 ml beaker
glass stirring rod
disposable cup
Hot plate
10 ml graduated cylinder
Hot mitts
De-ionized water
Polyvinyl alcohol (PVA) the powder form
Saturated Borax solution (add enough borax to water so that it turns cloudy. You can put it on a magnetic stirrer to keep the particles suspended)

  1. Mass out 2.00 grams of the PVA. Set aside.
  2. Pour 50 ml of deionized water into the beaker. Insert the thermometer into the beaker and place the beaker on the hot plate. Heat gradually to about 90 degrees. Do not let it boil rapidly or you will lose too much water and your slime will be stiff.
  3. SLOWLY sprinkle in the 2.00 grams of PVA and stir constantly with your glass stirring rod. You will know if you are going too fast if there is a glob of material at the bottom of your stirring rod.
  4. After you have completely stirred in all 2.00 grams of PVA, turn off the hot plate and keep your beaker on the hot plate so it doesn’t cool off. 
  5. Get 5.0 ml of Borax solution.
  6. Take your disposable cup and simultaneously pour the PVA solution from the beaker (use hot Mitts!!) and the Borax solution together in the disposable cup (NOT the beaker) and stir.
  7. A gel-like substance (SLIME) should form immediately. If it doesn’t, keep stirring. Sometimes when the solutions get too hot it takes a while to get the slime to form.
  8. Add food coloring to make really gross slime.
Now, why is this a Non-Newtonion Substance? Because it behaves differently depending on how gently or strongly you stir or pull it. If you pull it slowly, it will stretch and ooze sort of like a liquid. If you pull it apart quickly, it will stiffen up and break cleanly in two as if it were a solid. You’ve heard of quicksand, right? It is also a non-Newtonian substance. Maybe you’ve seen movies where someone is trapped in quicksand and cant’ get out. The harder they thrash around to get out, the worse it is for them. Can you explain why? Think about your slime. The harder you force it, the more rigid it becomes, so the person gets even more stuck in the quicksand. To save themselves, they should move very slowly to get out so they quicksand would behave more like a liquid and not resist the person as much. Ketchup is another non-Newtonian substance, but it behaves in the opposite way. Glass bottles of ketchup used to be common, but they are probably only seen now in restaurants. Ever try to get the ketchup out of this kind of container? It flows better with more agitation! You may have had to shake the bottle pretty violently several times before it actually starts to flow out of the bottle.
Easy Slime Alternative Lab
1 cup of cornstarch
1/2 cup of water
food coloring
1. Put cornstarch in bowl
2. Slowly and with stirring (hands are fine) add the water.
3. Add food coloring as desired
4. Test your cornstarch slime by hitting hard, then softly. Try to stir it quickly, then very slowly and gently. Note observations and have fun!
VI.  Desalinization by Distillation
This re-printed lab from Physical Science: Concepts in Action is courtesy of Pearson Prentice Hall. It instructs students in the concepts involved in the distillation process. They use both an active process (boiling) to distill salt water, and a passive process (solar evaporation). Click here to print a copy of the teacher’s version of this lab. Click here for more information on Prentice Hall’s Physical Science: Concepts in Action textbook and resources.

Hands On Science Activities Part 3

Pop-can calorimeter

“Counting calories in the classroom”

First of all, it’s “pop” here, not “soda”, “Coke”, or “sodee”.  This may not be true in your area. Check the indigenous population. Take an empty can. It needs to still have its ring. Fill the can about 1/3 full with water. Pull the ring up so that it is vertical. Put a glass rod or similar through the ring and suspend the can through the ring of a ring stand. Be careful so that the rod does not slip off. Use tongs or forceps to hold a small piece of potato chip.  Light the chip with a match and hold it under the can. The water heats up. Use this equation to find the amount of energy stored in the chip (or pretzel, peanut, whatever).

Heat gained or lost = (specific heat of substance) * (mass of substance in grams) * (temperature difference)

(Remember, the substance of interest here is the water because its temperature is easy to measure.) The specific heat of water is 1.00 cal/g*degree. Heat energy is measured in calories.  Remember that a food calorie (or Calorie) is 1,000 calories (the amount of heat energy required to raise the temp. of 1 g of water 1 deg. C).

Experiment idea: Try regular versus baked chips.

Here’s a photo from our calorimeter experiment; burning up Lays chips. We used the EasySense Datalogger (discontinued) to keep track of the temperature changes over time.


Energy conservation in a “Stopped Pendulum.”

Here’s an interesting demo to show potential energy and the conservation of energy. Make a pendulum by hanging a mass from a string.  Put a barrier in the path of the string, as shown in the diagram.  The pendulum swings down and the string effectively becomes shorter as it strikes the barrier.  How far up will the mass swing?  Lower, higher, or the same level as before? The mass still rises to about the same height. Put a board behind the pendulum and mark the starting height, or put it in front of a chalkboard, so that you can show that the heights are about equal.

TRY IT! This complete demonstration lesson is available here.

Newtonian Demonstrator - Newton's Cradle

In Stock SKU: P1-6001

Generators and Motors

Ever wish you could bottle the energy young children bring to school?  Now you can use it to help them learn!

Generators and electric motors; there are basically the same thing. But sometimes this simple concept is lost on some students. The Genecon hand-generator can be used to study all kinds of energy issues. Supplying up to 5 volts, it acts as a generator when you turn the handle to create an electric current, and as a motor when you supply current that causes the handle to turn.

The most obvious experiment is how much mechanical energy does it take to power light bulbs.  Students can really feel the increase in their effort as you add more light bulbs to the circuit.  Even young children can make the connection between more light bulbs and needing more energy to power them.

How about using it to power a Constant Velocity Car, instead of batteries?  (Or use the car to power the Genecon!)  How about using the Thermoelectric Device and turning the electricity generated by the Genecon into heat?  or cold?  (OK, removing heat).  Use it instead of batteries to power all kinds of 4-6V devices.  Figure out how efficient the Genecon is by powering one by another.  Count the number of times you turn the handle and how many turns you get from the other handle.  Challenge students to figure out where the lost energy went.

Genecon Hand Crank Generator

In Stock SKU: P6-2631

Constant Velocity Car (Carts)

In Stock SKU: 44-1090

Thermoelectric Device

In Stock SKU: P3-2600

Watts Up?

… the sky.  No, really. Waaaaaaatts Uuuuuuup?

Ok, the “Watts Up?” is really a meter.  You plug it into a standard wall outlet and plug an electrical device into the meter.  The meter will show you how much energy you are using and how much you have used since it was plugged in.  You can even set the rate so you know how much you would pay for that much electricity.

This is a great tool for units on energy conservation.  Compare different appliances.  Compare different sized light bulbs, or incandescent vs. Fluorescent vs. Halogen bulbs.  Challenge students to identify which device in a group uses the most electricity, and test their hypotheses.

Watts Up PRO

In Stock SKU: P6-8055

Sugar and Strength

Sugar – is there a link to strength?

Sugar provides energy. Does it also increase strength?

Have the students perform a strength test. It can be arm wrestling, lifting weights, or something else that makes them use almost all of their strength (at least in one limb). Have them take a pinch or cube of sugar and put it into their mouths. Try the strength test again. They will be amazed at how much less strength they have.  This works especially well when they are testing themselves near their limit. For instance, I did this with a student where we were dead even in arm wrestling. He took the sugar saying, “Now I should win quickly.” When I won quickly instead, he was shocked.

OK, but what’s happening? That’s not entirely clear. The effect is too quick and too drastic to be from digesting the sugar. What seems to be happening is that there is a feedback somehow to the brain. You taste sweet and the brain responds. Challenge the students to explain what they’ve observed. Consider working with a biology teacher to help students make the connection between physical and life sciences. Note: this experiment does not work right after lunch or after the students have had something sweet.