The Element Sulfur
Sulfur is a highly versatile element that is essential to life. This brittle, nonmetallic element belongs to Group 16 of the periodic table, along with Oxygen. Its characteristics include being odorless and having a distinctive pale yellow color.
Cool Facts About Sulfur
- Sulfur is the tenth most abundant element in the universe and the ninth most abundant element on Earth.
- There are at least twenty-two identified sulfur allotropes.
- China is the world’s largest sulfur producer. In 2019, the country produced 17.4 million tons.
- In its natural state, sulfur is known as brimstone.
- Sulfur is Latin for “burning stone”
- Planet Jupiter has a yellow moon, which is covered in sulfur and sulfur compounds.
- Fool’s gold (Pyrite), iron disulfide, contains sulfur in the rare -1 oxidation state.
- Joseph Priestly discovered sulfur dioxide by burning sulfur in 1774, calling it “vitriolic air.”
- Sulfur is spelled sulphur in the UK and Australia, but sulfur is the official IUPAC spelling.
- Pure sulfur is odorless, but its compounds can smell quite bad
- Hydrogen sulfide is extremely toxic, it can deaden your sense of smell and can cause death
Sulfur in the Periodic Table
Sulfur, with the atomic symbol S and atomic number 16, lies in group 16 of the periodic table. Group 16 elements are sometimes known as chalcogens. It lies to the right of the element phosphorus, and to the left of chlorine. Sulfur is below oxygen and above selenium, and shares some chemical properties with both. Just like selenium forms selenates and selenites, sulfur forms sulfates and sulfites.
Sulfur, a nonmetal, has 23 isotopes, 4 of which are stable and found in nature. Sulfur can be found in its elemental form, as bright yellow elemental sulfur. The sulfur atom has electron configuration [Ne]3s23p4 or 1s22s22p63s23p4 (full). Sulfur’s electronegativity is 2.58, similar to carbon.
Chemical & Physical Properties of Sulfur
Sulfur is a pretty, pale yellow color. It is a tasteless, odorless, and brittle non-metallic element. Additionally, this solid is insoluble in water and has poor thermal and electrical conductivity. Sulfur melts at a temperature of about 115.21˚C (239.38˚F). At temperatures of about 190˚C (374˚F) the runny, yellow liquid, turns thick and brown. As the temperature increases further, the liquid turns more viscous- until 300˚C (572˚F). At this point, the thick liquid becomes runny again.
Allotropes of Sulfur
An allotrope is a different structural form of an element resulting from different atomic arrangements. Sulfur allotropes are divided into two categories: crystalline and amorphous.
The two most important crystalline allotropes are the rhombic (octahedral, or α-sulfur) and monoclinic (β-sulfur) structures. Monoclinic sulfurs are amber-colored, transparent crystals, arranged in S8 rings. To make monoclinic crystals, crystallization occurs at a temperature of 100˚C. First, sulfur melts into a liquid. Then, it cools slowly. During the cooling stage, a solid crust forms on the surface. Next, the crust is punctured and molten sulfur is poured out. Once the crust is removed, monoclinic crystals should be revealed. These crystals have a density of 1.98g/cm3, which is less than elemental sulfur. It is important to note that these crystals only exist at temperatures above 96 ̊C. Below this temperature, the allotrope is unstable and converts to rhombic sulfur.
The rhombic structure is the most thermodynamically stable of all the sulfur allotropes. It contains S8 molecules that are aligned in symmetrical eight-membered rings. These S8 rings are packed differently than the monoclinic allotrope, resulting in different shapes. With a density of 2.08g/cm3, it weighs about the same as elemental sulfur. Additionally, these allotropes are insoluble in water but will dissociate in carbon disulfide solutions. To prepare rhombic sulfur, powdered sulfur is dissolved in carbon disulfide solution at room temperature. The solid that forms is filtered out and placed in a beaker and covered. Once the carbon disulfide evaporates, yellow, translucent rhombic crystals will remain.
Amorphous means to have no shape. Elements such as antimony and boron are also amorphous. When sulfur is heated to temperatures above 120˚C, it melts into a liquid. If heated more, about 200˚C, the liquid turns dark and thickens. This happens because high temperatures force the S8 rings of rhombic sulfur to break open. This results in multiple chain structures with exposed atoms at the end. These atoms possess electrons that absorb light well, explaining the change in color. Furthermore, multiple chains link to each other, forming a coil, increasing viscosity.
Colloidal allotropes and milk of sulfur exist in the amorphous category. Colloidal sulfur comes from processing hydrogen sulfide through a cooled solution of sulfur dioxide in water. The function of this allotrope is to act as a solvent in carbon disulfide and is utilized in medicine. Milk of sulfur is a product of the reaction between hydrochloric acid and ammonium sulfide. This allotrope has a white color and is non-crystalline. When heated, the white transforms into the known yellow color of sulfur.
Sulfur in Nature
Sulfur is extremely abundant in nature. Deposits of sulfur occur around hot springs and in volcanic regions. At these sources, hot gas expels and sulfur sublimes. Minerals including iron pyrites, galena, cinnabar, and Epsom salts contain the element as well. However, the greatest source of sulfur is evaporite minerals, such as gypsum and anhydrite. They are called evaporite minerals because they form when seawater evaporates. When deep in the ground, bacteria alter these minerals, forming sulfur.
Sulfur’s Applications in Today’s World
About nine million tons of sulfur are produced every year in the United States alone, and with how widespread its usage is, it is obvious to see why so much is needed. Pure elemental sulfur is used to make paper and is an ingredient found in insecticides and fungicides. Additionally, products such as gunpowder, matches, and fireworks rely on the element as well.
Sulfur is commonly in the form of sulfuric acid, which is the number one chemical in the industrialized world. Its greatest use is for fertilizers. Other products such as lead-acid batteries, pigments, detergents, and sheet metals contain the compound as well. Sulfur dioxide, another compound of sulfur, is a bleaching agent and disinfectant.
The human body requires sulfur for a few reasons. First, sulfur is essential to the building and fixing of DNA. It also protects cells from damage, which further helps to prevent diseases, including cancer. Additional ways sulfur keeps us healthy include:
- Keratin contains sulfur. It is a protein that promotes healthy hair and nails.
- Insulin levels are partially regulated by sulfur.
- Healthy joints need sulfur.
- Sulfur is involved in the synthesis of collagen, which is a protein that helps maintain healthy skin.
- Sulfur is a component of cysteine and methionine, which are amino acids.
History of Sulfur
Have you ever wondered about what pigment was used for cave paintings? The answer is sulfur. This element is ancient and has been used since prehistoric times. It was always recognized as a compound instead of its own element. It was not until 1777 that French chemist Antoine Lavoisier isolated sulfur. However, some still considered the element to be a compound of hydrogen and oxygen. In 1809, Louis-Joseph Gay Lussac and Louis-Jaques Thénard (who were also involved in the discovery of boron) proved the elemental nature of sulfur.
Sulfur is an extremely reactive element that reacts with all metals (excluding gold and platinum) to form sulfides. It is similar to iodine in that respect. Sulfides are inorganic, ionic compounds that are considered to be the salts of hydrogen sulfide. It is insoluble in water and is a good electrical insulator.
Excluding noble gases, the element sulfur forms stable compounds with all the elements, usually forming polycationic compounds.
Sulfides are anions/polyanions of the element sulfur. They generally exhibit basic behavior and therefore, are not that reactive towards acids. Instead, they react well with oxidizing agents; They will readily oxidize at room temperature in moist air or at elevated temperatures in dry air. For the most part, sulfides get hydrolyzed by water. Only the sulfides that react with either alkali metals or alkaline earth metals are considered to be soluble in water; Sulfides that react with copper and zinc are the least soluble. These compounds have diverse applications, ranging from lithium batteries, pigments, ceramic coding, lubricants, and more.
The classifications of sulfides are inorganic, organic, or phosphine.
Inorganic sulfides contain a negatively charged sulfide. They are usually found in sediments and are the product of a sulfur atom reacting with a metal. A well-known inorganic sulfide compound is hydrogen sulfide, H2S. Among the inorganic sulfides, hydrogen sulfide is the most dangerous in terms of it being extremely flammable and toxic. However, other inorganic sulfides are also capable of igniting when heated and exposed to moisture. When exposed to H2S gas, it can cause central nervous system (CNS) and respiratory depression.
More mild effects include headaches, and eye and skin irritation. What is even more frightening is that there is currently no proven medication to treat H2S poisoning. Another characteristic that distinguishes H2S from other inorganic sulfides is that it is volatile- it evaporates easily. Upon vaporizing, an unpleasant, rotten egg smell is released. The other compounds are generally solid and are not evaporative, and therefore do not smell bad (thankfully).
Organic sulfides (R-S-R’), sometimes called thioethers, are analogs of ethers (R-O-R’). In these compounds, sulfur is covalently bound to two organic groups, which are either aliphatic, aromatic, unsaturated or a combination of the three. Organic sulfides can be found abundantly in nature- even in garlic (allyl sulfides are found here)! When an organic sulfide reacts with a strong reducing agent, it will usually result in the release of hydrogen gas. While hydrogen gas is not considered to be toxic, it is flammable when combined with oxygen in the air. Organic sulfides are mainly used as additives in petroleum and pesticides, and in the synthesis of chemicals, such as dimethyl sulfone (as shown in the diagram).
A Thiol is a type of organic functional group which is the sulfur analog of alcohols. They have the structure R-SH. As with many sulfur compounds, thiols often have distinct odors. They occur in nature as a part of many biological systems. 2-butanethiol is also a component in skunks’ spray.
Phosphine sulfides are composed of a sulfur atom bonded to a phosphorus atom. They are derivatives of organophosphines. They are pentavalent, meaning they have five valence electrons. Additionally, they are stable, soluble, and have low polarity, making them great intermediates in the chemical reactions of organophosphines.
Sulfur oxides are the product of burning sulfur and oxygen. Common oxides include sulfuric acid (H2SO4), sulfur dioxide (SO2), and sulfur trioxide (SO3). While all sulfur oxides are damaging to the environment and our health, sulfur dioxide is the greatest attribute. This is because other sulfur oxides are produced from sulfur dioxide- While there are other ways for the other oxides to be produced, sulfur dioxide is mainly responsible for their existence.
Sulfur dioxide is a colorless gas that is found in the lower atmosphere. It has an unpleasant odor and taste, that can be detected at 1000 to 3000 µg/m3. Human activity accounts for the greatest amount of sulfur dioxide emissions; Burning fossil fuels releases the greatest amount of sulfur dioxide, but it also comes from smelting sulfide ores. While natural sources like volcanoes, and biological decay, also emit sulfur dioxide, it only accounts for a small amount. This gas is horrible for both the environment and people. It harms nature by damaging plants and trees, which can negatively affect the growth of new plants.
Sulfur dioxide affects the lungs’ function, which can lead to respiratory problems. It also irritates the eyes, nose, and throat. People with respiratory diseases, like asthma, are even more prone to the dangers of sulfur dioxide. Additionally, when sulfur oxides react with other compounds in the atmosphere, it forms small particles that can penetrate into the lungs. These particles also contribute to the haze that we sometimes see outside.
Sulfur trioxide forms when sulfur dioxide and oxygen combine. It can exist as a colorless liquid, ice-like crystal, or gas. As a gas, it is white, foggy, and smelly. Even in small quantities, the smell can be quite pungent. Sulfur trioxide is highly reactive, making it ideal for the synthesis of other chemicals, like sulfuric acid, and for producing explosives.
Sulfur trioxide combines with the water and oxygen in the air to form sulfuric acid. This is a colorless liquid that is strongly acidic, corrosive, and heavy, with a density of 1.83g/cm3. When sulfuric acid dissolves in the water vapor in the air, it turns into acid rain. Sulfuric acid readily dissociates in water, forming H+ and HSO4–. This ion is considered to be a weak acid. It can dissociate again to form H+ and SO42-. When the concentration of H+ ions increases, the pH of the rain decreases, hence the name acid rain. Acid rain is harmful to plants and can contaminate the soil. It also disturbs the pH levels of lakes and streams, which endangers aquatic wildlife. Additionally, it damages the surfaces of buildings, statues, and monuments and can cause paint to peel.
Isolation of Sulfur
Herman Frasch is credited for inventing the Frasch process, which extracts 99% pure sulfur. Its purpose is to mine elemental sulfur from deep wells within salt domes. The process works by first heating water to a temperature of 170˚C (340˚F) and then forcing it into the wells. Sulfur has a low melting point of 115.21°C (239.38°F) and will melt from the hot water. Then, the element rises to the surface due to compressed air. Today, the Frasch process is used in Poland and Russia.
There are other methods of obtaining pure sulfur as well. It is more common to obtain sulfur through purifying sulfur-rich petroleum crudes and gas sources. This method works by heating the sources at a boiling point, which will result in sulfur separating.
Sulfur Oxidation States
Sulfur exists in oxidation states ranging from -2 to +6. Depending on how many bonds sulfur has, it will have a different oxidation state. For example, in sulfuric acid (H2SO4), sulfur has an oxidation state of +6. To obtain a stable solution of polycations, sulfur gets oxidized with SO2 in an oleum solution. When a sulfur ion and a mild oxidizing agent, such as S2O6F2, AsF5, or SbF5 mix in a strong acid solution, they form colored polycations, like S162+, S82+, and S42+. S162+ is red, S82 is dark blue, and S42 is yellow. These polycations all have oxidation states of +1 or less.
Physical Properties of Sulfur
- Melting point: 388.36 K; 115.21°C; 239.38°F
- Boiling point: 717.76 K; 444.61°C; 832.3°F
- Density: 2.07 g/cm3
- Atomic weight: 32.06
- Atomic number: 16
- Electronegativity: 2.58
- Classification: Nonmetal
- Natural abundance in the Earth’s crust: 0.03%
- Electron shell configuration: [Ne] 3s23p4
- Isotopes: Sulfur has four stable isotopes, which are S-32, S-33, S-34, and S-36
- Found naturally in the minerals: Sulfate and sulfide minerals, such as pyrite, galena, bornite, realgar, and cobaltite
- Toxicity: Low toxicity
Where can I buy Sulfur Element?
Sulfur powder can be purchased from Amazon, eBay, and specialty stores.
Sulfur crystals can also be purchased from Amazon, eBay, and specialty stores.