When introducing the elements, it is always more fun for the students if they have the opportunity to examine a few rather than just hearing about them. Sulfur presents an excellent opportunity, as it has a fairly unique set of properties, without being terribly hazardous. The following activities will explore a number of the chemical and physical properties of sulfur.
(Safety note – rare individuals are allergic to sulfur, check for known sulfur allergies, and be alert to symptoms when sulfur is exposed to the classroom. Working in a fume hood at all times minimizes this risk.)
#1 – Color and Form
Perhaps not the most exciting observation, but considering how few elements have bright colors to boast of, it could be worse. They should study the texture as well, as it leads in to #2.
#2 – Metal or Nonmetal?
Students should test a few properties to be sure. Hopefully they’ve already studied this, and know that metals are typically shiny, malleable, ductile, and good conductors, whereas nonmetals tend to be dull, brittle, and bad conductors. Have a conductivity probe handy for the last one. Better yet, let the students make one using a simple open circuit with a light/buzzer and a battery.
#3 – Solubility
Test the solubility of sulfur by sprinkling a small pinch of powdered sulfur into about 50 mL of each solvent. Try water (polar), ethanol or ether (somewhat polar), and benzene or hexane (non-polar). Ether, benzene, and hexane should only be used if you have a fume hood.
The results of this test should give an indication as to whether elemental sulfur is polar or non-polar.
If you wish to, you can also demonstrate that sulfur is quite soluble in carbon disulfide, but since this is more toxic, don’t let the students be involved at this step, and stay in the fume hood.
#4 – Combustibility
Sulfur will burn, producing sulfur dioxide gas. If inhaled, it dissolves in the moist linings of nose, throat, and lungs, forming sulfurous acid and causing a great deal of pain and coughing. To avoid this, conduct all combustibility experiments in the hood.
If equipment is limited, the simplest approach is to ignite a small spoonful of sulfur over a Bunsen burner. This produces a cloud of smoke, dumping sulfur dioxide out into the air. While presenting an opportunity to discuss acid rain, this is not the best way to go.
A better approach is to ignite the sulfur in a container that is sealed, save for one opening with a piece of tubing that leads out and down into a beaker of water. The water should have an acid-base indicator added to it, and just enough base added to shift the color to the basic color. Phenolphthalein is a good choice, starting off pink, and shifting to clear when the sulfur dioxide gas acidifies the water. Note that intense heating will ignite sulfur – a flame is not necessary.
This gives a visual demonstration of how burning sulfur leads to acid rain, as the sulfur dioxide gas dissolves in the water in the air. Discuss how sulfur is a component of coal, and how burning coal releases sulfur dioxide in the same way. You can also talk about methods used to reduce sulfur dioxide emissions. The subject also touches on history, since gunpowder traditionally contained sulfur. Historic battles (like the Revolutionary and Civil Wars in the United States) were fought using this gunpowder, releasing clouds of sulfur dioxide that rolled over the battlefield, where it would burn eyes, noses, throats, and lungs. If hell is fire and brimstone (the ancient name for sulfur), it’s no wonder they say that “war is hell”.
#5 – Melting Point and Allotropes
Sulfur doesn’t melt until a little bit above the boiling point of water, so an oil bath is best for this experiment. (You can use a melting-point apparatus if you want for a more precise melting point determination, but it won’t give you enough material for the rest.) A beaker of silicone oil on a hot plate with a stir bar will make an excellent bath. Fill a small test tube about one third to one half full, and hold it (with a test tube holder) in the oil bath. Monitor the temperature of the bath, heat to 100°C, and then slowly increase the temperature until the sulfur melts. (Record both the temperature and the color.) Heat for a couple degrees more, then quickly dump the contents of the test tube into a small beaker of water.
Pour off the water, dry the sulfur, and examine its color and form once more. While the original sulfur was an elemental form containing eight-atom rings (or crowns), the heating step caused those rings to break open, and the rapid cooling by plunging the molten sulfur into water prevented the rings from reforming. Instead, the sulfur is now a chain form of sulfur, called plastic, or amorphous, sulfur. These two different forms of the same element are called allotropes.
