Element of the Month - Sulfur

"In honor of the International Year of the Periodic Table this series of articles details the Element of the Month project developed by Stephen W. Wright (SWW), Associate Research Fellow at Pfizer Inc., and Marsha R. Folger (MRF), chemistry teacher (now retired) at Lyme – Old Lyme High School in Connecticut. Read The Element of the Month - An Introduction for an overview of the project and links to the other articles in the series." - Editor

The fourth element highlighted in our Element of the Month program is sulfur. For us, this meeting would occur in December. Sulfur plays an essential role in everyday life but it is likely that the students will have little familiarity with sulfur and its compounds. The inclusion of sulfur in the Element of the Month program allows them to make connections between sulfur and its many roles. The December class schedule is of course compressed due to the holidays, and the sulfur demonstrations are designed to be relatively simple to arrange.

Figure 1: Sample of elemental sulfur

Occurrence in Nature

Students will probably be unaware that a vast amount of sulfur occurs in the form of sulfate ion in seawater. Sulfur also occurs in numerous sulfide and sulfate minerals, some of which are important ores. Sulfur can be found as the free element in certain situations, and some students may know that sulfur is frequently associated with volcanoes and hot springs. We have large samples, of 100 to 300 g mass, of pyrite (“fool’s gold”, FeS₂), galena (PbS), gypsum (CaSO₄), and elemental sulfur (see figure 1). These are contained in clear plastic jars and passed around the class for inspection. We note that sulfur is also found in coal and petroleum. Lastly, we explain that sulfur occurs in the amino acid cysteine, which plays a key role in maintaining the three dimensional structure of proteins, especially hair, horn, and claws.

Uses

The largest use of sulfur is in the manufacture of sulfuric acid, which students will realize is “battery acid”. They will respond that they know that sulfuric acid is “strong”, but they will not know that sulfuric acid is and always has been the largest production volume industrial chemical. Sulfur and its compounds are used in vulcanizing rubber and in the manufacture of a great many chemicals. Many students will know that sulfur is a component of the old fashioned explosive black powder, and in matches and pyrotechnics. Most detergents are compounds of sulfur. On the lecture table, we display various products that contain sulfur, including a small tire such as a bicycle tire, a battery from a lawn tractor or other small vehicle, detergent, Plaster of Paris, a scrap of gypsum wallboard, Epsom salt, and match box.

Physical Properties

We take a few minutes to note that sulfur is bright yellow, odorless, insoluble in water, brittle, and does not conduct electricity. It melts at 115 °C, just above the boiling point of water.

Figure 2: Reaction of zinc and sulfur. A) Mixing zinc and sulfur, B & C) Two views of heated rod and zinc/sulfur mixture on heat resistant mat, D) Ignition of mixture, E & F) After ignition. Derived from Jerrold J. Jacobsen and John W. Moore. Chemistry Comes Alive: Vol. 3: Abstract of Special Issue 23 on CD-ROM, Journal of Chemical Education 1999 76 (9), 1311. DOI: 10.1021/ed076p1311.

Chemical Properties

We first explain that sulfur chemistry is noted for the various oxidation states that sulfur can take. On the board, we show the formulas and oxidation states of sulfur in sulfur (S, 0), sulfide (S^2-, -2), sulfite (SO₃^2-, +4), and sulfate (SO₄^2-, +6). We observe that sulfur can combine directly with some metals such as iron and zinc to form sulfides. We demonstrate the very exothermic reaction between zinc and sulfur to form zinc sulfide by mixing 4 grams of gray powdered zinc metal and two grams of yellow powdered sulfur.¹ We place the mixture on a heat resistant mat in the fume hood and ignite the mixture using a gas burner or heated rod (figure 2). After the reaction has cooled, we remove the mat and show the class the white product. We note that many of the transition metals form extremely insoluble sulfides and are often mined as their sulfide ores.

Figure 3: Precipitation of the sulfides of Cd, Pb, Mn and Zn in sodium sulfide solution

We show the precipitation of the sulfides of cadmium, lead, manganese, and zinc by placing aqueous solutions of these metals in large test tubes in a test tube rack and treating the solutions with a little dilute aqueous sodium sulfide solution (see figure 3).² Colorful sulfide minerals such as cinnabar (HgS, red), orpiment (yellow, As₂S₃), and stibnite (black, Sb₂S₃) were used as pigments in ancient times. We explain that sulfur also forms a sulfide compound with hydrogen, and ask the class if they know what that substance is. We explain that hydrogen sulfide is a colorless, flammable, poisonous gas that has a very familiar, rotten egg odor. As an aside, we note that burning rubber stinks because of organic sulfur compounds used to formulate the rubber. Some students will know that hair perm solutions smell, and we explain that this is also from sulfur compounds that are present. As a general rule, sulfides smell, while oxidized sulfur compounds generally are much less odorous.

Figure 4: A) Sulfur powder in spoon before reaction, B) Blue flame of burning sulfur in darkened room, C) Sulfur dioxide fog after addition of water and shaking of the flask, D) After addition of water and universal indicator.

Sulfur forms compounds with oxygen too. For example, it burns in air to form sulfur dioxide. Sulfur dioxide in the environment comes from volcanoes, from burning coal or oil that contains sulfur, and from “roasting” metal sulfide ores. Sulfur dioxide is a colorless, toxic, dense gas, one that does not burn or support combustion. It has a choking, familiar smell that many associate with the smell of burning matches. We burn a bit of sulfur using a long handled deflagrating spoon in a 1 L flask in the darkened classroom, and we note that the sulfur burns with a clear blue flame (see figure 4). After a short while, we extinguish the sulfur and add a small amount of water, about 50 mL, to the flask. We stopper the flask and shake the flask, then gently release the stopper. An audible sound of air rushing into the flask demonstrates that sulfur dioxide is extremely water soluble.³ We add some universal pH indicator solution to the flask and show that the sulfur dioxide solution gives an acidic reaction with a pH indicator. Students will usually have heard of acid rain and this is an opportunity to make that connection. We will usually add some indicator to a large test tube of water as a control, and to a test tube of a dilute acid as well. If a bell jar or similar large container is available, one can show how sulfur dioxide may be used as a bleach. A red or violet carnation is placed under a bell jar with some burning sulfur in the fume hood (see figure 5).⁴ By the end of the class period, the carnation will have lost most of its color. We explain that the acidic solution formed from sulfur dioxide in water can be neutralized with a base, and the resulting salts are known as sulfites. These compounds are reducing agents and are frequently used as preservatives. For example, produce and wines may be preserved with sulfites. As a demonstration, we show the class two petri dishes, each containing a slice of very ripe apple or pear. One of the slices is untreated and brown, while the second half has been treated with a few drops of a sodium bisulfite solution and appears fresh. We will usually note that sulfites are used to manufacture cheap paper, but that paper made with sulfites will gradually deteriorate with age due to the acidic nature of sulfites and sulfur dioxide. Lastly we show the reduction of permanganate ion by sodium bisulfite solution by pouring some dilute potassium permanganate into a sodium bisulfite solution.⁵ 

Figure 5: A) A tired carnation, bell jar and dish with powdered sulfur,  B) Carnation and dish together under bell jar ready for ignition, C) After 30 minutes, D) Carnation is mostly bleached except wilted edges.

Sulfur’s most common oxidation state is +6, as found in the sulfate ion. Sulfate ion is found in nature as calcium sulfate (the mineral gypsum) and magnesium sulfate (the mineral epsomite, or Epsom salt). We note that gypsum is used to make plaster, drywall, and cement. We explain that sulfuric acid is made from SO₂ by oxidation to SO₃ followed by the reaction of SO₃ with water. Sulfuric acid is indeed a very strong acid, because it completely dissociates in water, with two protons per molecule. We make it clear that concentrated sulfuric acid is pure H₂SO₄, unlike many other concentrated acids found in the laboratory (such as HCl and HNO₃) which are actually solutions in water. We show the heat of dilution of sulfuric acid with water in a Pyrex flask, slowly adding 20 mL of concentrated sulfuric acid to 200 mL of water to produce a 20 °C temperature increase.⁶

Video 1: Dehydration of sucrose by sulfuric acid (accessed May 2019 subscription required) Derived from Gary Trammell, Jerrold J. Jacobsen, Kristin Johnson, and John W. Moore. Chemistry Comes Alive! Volume 5: Abstract of Special Issue 29, a CD-ROM for Organic and Biochemistry, Journal of Chemical Education 2001 78 (3), 423. DOI: 10.1021/ed078p423.

Next we demonstrate the charring of sugar by sulfuric acid by adding 35 mL of concentrated sulfuric acid to a 250 mL beaker containing a mixture of 35 g of granulated sugar and 35 g of confectioners’ sugar (see video 1).⁷ This is stirred with a glass stirring rod for a few seconds and then placed in the fume hood.

Figure 6: A) Sulfuric acid was used to write H₂SO₄, B) Use of a heat gun reveals the message

Lastly we write a “secret” message on a sheet of paper (see figure 6) using 3 M sulfuric acid and expose the message by heating the paper using a heat gun, while taking the opportunity to stress the need for laboratory safety and awareness around corrosive chemicals by reciting the following poem:⁸

Little Willie took a drink,

And lived to tell no more

For what he thought was H₂O

Was H₂SO₄

For cleaning up, we neutralize the diluted sulfuric acid solution resulting from the heat of dilution demonstration using magnesium hydroxide powder to permit disposal down the sink. Each 20 mL of concentrated sulfuric acid requires approximately 21 grams of magnesium hydroxide to fully neutralize the acid. The advantages of magnesium hydroxide powder are that almost every high school laboratory has this compound and if an excess is used, the pH of the solution remains neutral.

References and Notes

1. An excellent discussion of the presentation of this demonstration may be found in: Shakashiri, Bassam A. Chemical Demonstrations: A Handbook for Teachers, Vol. 1; University of Wisconsin Press: Madison, WI, 1983; pp 53-54.
2. See Abstract 18-24 in Alyea, Hubert N. Tested Demonstrations in Chemistry, 6th ed.; Journal of Chemical Education: Easton, PA: 1965; pp 40.
3. (a) See Abstract 19-16 in Alyea, Hubert N. Tested Demonstrations in Chemistry, 6th ed.; Journal of Chemical Education: Easton, PA: 1965; pp 42; (b) Shakashiri, Bassam A. Chemical Demonstrations: A Handbook for Teachers, Vol. 2; University of Wisconsin Press: Madison, WI, 1985; pp 184-189.
4. See Abstract 19-15 in Alyea, Hubert N. Tested Demonstrations in Chemistry, 6th ed.; Journal of Chemical Education: Easton, PA: 1965; pp 42. Bell jars are expensive and it is possible to use an inexpensive substitute. For example, the top may be cut off a large plastic water bottle and the bottle inverted over the carnation and burning sulfur. A disposable container such as a metal jar lid is used to hold the burning sulfur.
5. See Abstract 19-17 in Alyea, Hubert N. Tested Demonstrations in Chemistry, 6th ed.; Journal of Chemical Education: Easton, PA: 1965; pp 42.
6. Laboratory safety instructions always call for concentrated acid to be added into water, and note that the water should never be added to the acid. The origin of this practice lies in the density and the heat of dilution of concentrated sulfuric acid. When poured into water, concentrated sulfuric will sink to the bottom of the water where the heat of dilution may be dissipated by the bulk of the water. By contrast, if water is added slowly to concentrated sulfuric acid, the water floats on the sulfuric and immediately becomes so hot as to boil, frequently spattering acid out of the container.
7. Summerlin, Lee R.; Borgford, Christie L.; Ealy, Julie B. Chemical Demonstrations: A Sourcebook for Teachers Volume 2, 2nd ed.; American Chemical Society: Washington, DC, 1988; pp 122. We use a mixture of confectioner’s sugar and granulated sugar to optimize the time required and the height of the carbon pillar. If it is desired to clean the beaker, the charred sugar residue should be removed from the beaker while still very warm and the beaker soaked in water overnight.
8. The author (SWW) first found this poem in a high school yearbook belonging to the Aurora, NE class of 1916.