How to blow out a light bulb
One of the many advantages of having students peruse sources of chemical demonstrations rather than my telling them which ones they should perform is that they occasionally come across demonstrations that I have not seen before. One example of this is when Stacey Wolfson found “How to Blow Out a Light Bulb” in A Demo A Day, Volume 2, B. Bilash II, G. Gross, & J. Koob, (1998), pages 208-209, Flinn Scientific, Inc. This demonstration was new to me. As I watched Stacey practice the demonstration, a multitude of chemical concepts came to mind — conductivity of solutions, electrolytes versus non-electrolytes, solubility, precipitation reactions, acid-base chemistry, and chemical equilibrium; even a little biology comes into play. And, what a simple way to visually illustrate all of these concepts!
In this demonstration a conductivity apparatus is inserted into a solution of calcium hydroxide; the lamp glows brightly. Using a straw, the demonstrator steadily blows into the solution; the lamp dims and eventually no longer glows. A white precipitate, calcium carbonate, forms during this process.
CO2(g) + Ca2+(aq) + 2OH-(aq) → CaCO3(s) + H2O(l)
As shown in the equation, the calcium hydroxide is dissociated into ions, and therefore the solution will conduct electricity; the solution is an electrolyte. Approximately four-tenths (0.4) of a percent of our exhaled breath is carbon dioxide, which reacts with the calcium hydroxide solution. Eventually, “all” the calcium and hydroxide ions are converted into insoluble calcium carbonate and water; both can be considered as non-electrolytes. Without “any” ions to carry the electrical current through the solution, the light bulb does not glow.
The reaction can be viewed as a series of steps.
CO2(g) ⇌ CO2(aq)
CO2(aq) + H2O(l) ⇌ H2CO3(aq)
H2CO3(aq) ⇌ H+(aq) + HCO3-(aq)
HCO3-(aq) ⇌ H+(aq) + CO32-(aq)
Ca2+(aq) + CO32-(aq) ⇌ CaCO3(s)
H+(aq) + OH-(aq) ⇌ H2O(l)
The equilibrium constants for the first four steps shown are small and get progressively smaller. The stage at which the carbonate ion actually reacts with the calcium ion is uncertain. In any case, the calcium ion binds with the carbonate to form insoluble calcium carbonate, a precipitation reaction, and the hydrogen ion binds with the hydroxide ion to form water, an acid-base reaction.
Even though the equilibrium constants for the first four reactions are small, the formation of calcium carbonate and water drives all of the reactions in the forward direction.
- Dissociation, solubility
- Electrolytic solutions vs non-electrolytic solutions
- Precipitates and the formation of precipitates
- Acidic nature of carbon dioxide
- Acid-base reactions
- Chemical equilibrium — Le Châtelier’s principle
- 1 gram* of solid calcium hydroxide
- ~ 50 mL of distilled/de-ionized water
- Glass stirring rod
- Filter paper
- 100-mL erlenmeyer flask with stopper
- 250-mL beaker
- Conductivity apparatus
- Ring stand
- Plastic straw
*In the original source, the amount of calcium hydroxide was given as 10 grams. This seemed excessive since the solubility of Ca(OH)2 is 0.17 g/mL at 20∘C. This would make a good Ksp question for students.
- Wear safety glasses and protective gloves
- If the conductivity apparatus has bare wires, use caution making sure not to come in contact with the wires; electrical shock hazard
- Avoid contact with, or inhalation, of the calcium hydroxide; skin irritant
- Make sure demonstration volunteers “blow” into the straw and not “suck” in — small children sometimes have a problem with this.
- The calcium hydroxide solution should be fresh. Prepare it the day of the performance of the demonstration. Add 1 gram of calcium hydroxide to about 50 mL of distilled/de-ionized water and stir thoroughly. Allow the solid to settle. Carefully filter the solution into the 100-mL Erlenmeyer flask. Stopper tightly.
- Transfer the filtered calcium hydroxide solution to the beaker.
- Make sure the conductivity apparatus is NOT plugged into the electrical outlet. Attach the conductivity apparatus to a ring stand.
- Insert the electrodes into the calcium hydroxide solution. Be sure they are not touching the sides or bottom of the beaker, or each other.
- Insert the plug into the outlet. The lamp should glow brightly.
- Using the straw, steadily blow into the solution being sure that you do not come in contact with the wire electrodes or inhale some of the solution. As you blow into the solution the lamp will dim and eventually go out. Calcium carbonate, a white precipitate, will form as you blow into the solution.
- Discard the straw into the trash.
- The remaining solution in the beaker can be filtered. The solid can be disposed of in the trash and the liquid in the sink. Be sure to consult your local community and school policies before doing this.
As previously stated, Stacey came across this demonstration, tried it out, and performed it in front of her fellow classmates. Hopefully, she will have the opportunity to present this activity at an outreach event this coming academic year.
The following are Stacey’s thoughts and comments about “How to blow out a light bulb”:
Staging a demonstration in front of classmates is one requirement in the chemistry outreach course. I wanted to perform an exciting yet different experiment. My quest began by searching through recommended demonstration books. A few experiments were interesting, but after conferring with Dr. Lyle, these were deemed either too difficult to perform or too similar to the current demonstrations at our outreach events. When I came across “How to blow out a light bulb” I saw that it fit my requirements, new — to me at least — and exciting. This demonstration was particularly intriguing. I had previously seen the conductivity test for solutions but not this spin on the demonstration. I did not imagine there was a way to extinguish the glow of the bulb. I was fascinated by this experiment and the simple explanation of the underlying science.
To enhance my presentation I practiced several times in front of Dr. Lyle and other classmates. I now understand the importance of feedback because each observer has a unique perspective of the demonstration. I included several of these recommendations which enhanced my presentation. Based on observing presentations of my classmates, I incorporated some of their ideas and techniques. For example, one student wrote chemical equations on note cards and posted these as she explained the chemistry of her demonstrations. The class remained attentive and seemed to have acquired an accurate picture of what was happening. I decided to use note cards as well.
I geared my presentation to a high-school aged audience. First, I explained that lime water is a solution of calcium hydroxide and showed how the light bulb shined brightly when the electrodes were inserted in the lime water. Then, with the help of volunteers from the audience, we proceeded to “blow out” the light bulb. I found that asking questions helped the students to maintain focus and, as Dr. Lyle has stated on many occasions, informed me of the level of understanding of the audience.
Seeing everyone enthusiastic, willing to answer questions and eager to participate, made performing this demonstration an enjoyable experience.
*Stacey Wolfson is a senior at Duke University majoring in Spanish with minors in chemistry and biology. She plans to pursue a career in medicine upon graduation.
**Kenneth Lyle is in the Department of Chemistry at Duke University.
The Powell Family Trust, the Duke-Durham Neighborhood Partnership, and Biogen Idec – Research Triangle Park, fund the Duke Chemistry Outreach Program.