CHEMISTRY OF SULFUR

BY:

I KADEK IRVAN ADISTHA PUTRA 

I KOMANG ADI WIRANATA 

CHEMISTRY EDUCATION DEPARTMENT

FACULTY OF MATHEMATICS AND NATURAL SCIENCES

GANESHA UNIVERSITY OF EDUCATION

2012

            Sulfur is well known by the peoples because it is used in people’s daily life. For matches, gunpowder, detergents, fireworks, batteries are some of sulfur usefulness. Not only its usefulness, but also it is needed to know or make the others   know, the danger of sulfur itself. The usefulness and the danger are absolutely related into the science of matter, chemistry. In chemistry term, sulfur can be analyzed in the periodic table of the elements. Sulfur is placed at group VI A, and at the third period. Started from periodic table, it can be known all about sulfur including its properties and its allotropes, ways in producing sulfur, the kind of sulfur compounds and ways in producing or preparing sulfur compounds. So, this essay is going to tell or explain sulfur including its histories, its properties, allotropes, compounds and ways in producing sulfur compounds.

Sulfur was firstly known in China, in a natural form that the chinese had called “brimstone” or shiliuhuang that was found in Hanzhong (Freshney, 2007). By the third century, chinese discovered that sulfur can be extracted from pyrite. A Song Dynasty of military treatise of 1044 AD described different formulas of chinese gunpowder, which is a mixture of potassium nitrate (KNO₃), carbon, sulfur. Early, alchemist gave sulfur its own alchemical symbol which was a triangle at the top of a cross. In the late 1770s, Antoine Lavoisier helped convince the scientific community that sulfur was an element and not a compound. In 1867, sulfur was discovered in underground deposits in Louissiana and Texas. The overlying layer of earth was quicksand, prohibiting ordinary mining operation. Therefore, the Frasch process was utilized to produce sulfur.

In modern era, its specifically known the properties of sulfur and its allotropes. Sulfur is classified into group VIA in periodic table or it is called chalcogen group. Sulfur is also classified into non-metallic group. Sulfur has melting point at 388,36 K or 115,21^o C. The melting point of S₈ is actually a decomposition point. Just after melting rings with an average of 13.8 sulfur atoms are formed and at higher temperature still larger rings form. The in the high viscosity region there are giant macromolecules that are probably chains with radical ends. At higher temperature, highly colored S₃ and S₄ molecules are present to the extent 1-3 % at the boiling point.. Then, sulfur also has boiling point at 717,8 K or 444,6^o C. Sulfur is normally seen in lemon yellow colour. The density of sulfur is 2,08 gr/cm³ (alpha), 1,96 gr/cm³ (beta) and 1,92 gr/cm³ (gama).

Sulfur has allotropes. At least, twenty four  sulfur allotropes has been identified. Two of the most common sulfur allotropes are Orthorhombic sulfur (α-S₈) and Cyclohexasulfur or Rhombohedral sulfur (S₆). The most stable allotrope form in room temperature is Orthorhombic sulfur (α-S₈). Eight sulfur atoms bond covalently crown like rings. The Cyclohexasulfur or Rhombohedral sulfur is the densest of the sulfur allotropes and forms air sensitive orange red crystals containing chair shape, six membered rings. The structure of Orthorhombic sulfur (α-S₈) and Cyclohexasulfur are shown in figure below.

(a)

(b)

 

	Figure 1. The structure of Orthorhombic (a) and Cyclohexasulfur (b). The picture of (a) is taken from Garry L Miessller et al., “Inorganic Chemistry”. The picture of (b) is taken from Bodie E. Douglas et al., “Structure and Chemistry of Crystaline Solid”.

 

Sulfur occurs widely in nature as the element, as H₂S and SO₂, in metal sulfide ores, and as sulfates [e.g., gypsum and anhydrite (CaSO₄), magnesium sulfate and so on](Cotton et al., 1995). Freshney (2007) mention sulfur can be found in pure form (near hot springs and in volcanic regions) and in ores like cinnabar (HgS), galena (PbS), alunite, barite (BaSO₄), sphalerite (ZnS), and stibnite (Sb₂S₃). The main concern is the existence of sulfur as element. Sulfur as element is found in Orthorhombic sulfur structure (α-S₈). Related with sulfur usefulness, it is made a system or prototype to produce sulfur from natural sulfur resources.

In 1904, Frasch was successful to develop a method to extract sulfur from natural sulfur resources. This method is known as Frasch method. In this process, metal pipe with diameters of 15 cm which is consisted of two smaller concentric pipes inside is buried until reach the sulfur layer. Extremely hot water vapor is pumped into the sulfur layer through the pipe until the sulfur melt. Then, high pressure air is pumped into the sulfur layer through small pipe until sulfur foam is formed and the sulfur foam is pumped up into surface. On the surface, liquid sulfur is gathered in together in large vessel and let them freeze, produce solid sulfur with purity of 99,5%. The frasch method is shown in picture below.

	Figure 2. The pump in Frasch method. The picture is taken from C. Chambers et al., “Modern Inorganic Chemistry”.

 

Sulfur has oxides, especially sulfur dioxide. The dioxides are commonly produced by burning elements in the air. So, sulfur dioxides is produced when sulfur is burned in the air. Sulfur dioxide is a gas with a pungent smell (Cotton et al., 1995). The molecule is angular. Sulfur dioxide has lone pairs and can act as base according to Lewis acid-base theory. The other sulfur oxides is sulfur trioxide. Sulfur trioxide is obtained by reaction of sulfur dioxide with O₂, a reaction that is thermodynamically very favorable but extremely slow in the absence of a catalyst such as platinum sponge, V₂O₅, on NO (Cotton et al., 1995). The sulfur trioxide molecule, in the gas phase, has a planar, triangular structure involving both p-p an p-d S-O bonding and forms polymers in the solid state.

The other compound of sulfur is Sulfuric Acid. Sulfuric Acid, H₂SO₄, corrosive, oily, colorless liquid, with a specific gravity of 1.85. It is formed from hydrogen, sulfur, and oxygen. Sulfuric acid melts at 10.36°C (50.6°F), boils at 340°C (644°F), and is soluble in all proportions in water. When sulfuric acid is mixed with water, considerable heat is released. Unless the mixture is well stirred, the added water may be heated beyond its boiling point and the sudden formation of steam may blow the acid out of its container. The concentrated acid destroys skin and flesh, and can cause blindness if it gets into the eyes. The best treatment is to flush away the acid with large amounts of water. Despite the dangers created by careless handling, sulfuric acid has been commercially important for many years.

Sulfuric acid is a strong acid, that is, in aqueous solution it is largely changed to hydrogen ions (H⁺) and sulfate ions (SO₄^2-). Each molecule gives two H⁺ ions, thus sulfuric acid is dibasic. Dilute solutions of sulfuric acid show all the behavior characteristics of acids. They taste sour, conduct electricity, neutralize alkalies, and corrode active metals with formation of hydrogen gas. From sulfuric acid one can prepare both normal salts containing the sulfate group, SO₄, and acid salts containing the hydrogen sulfate group, HSO_4.

There are two major processes (lead chamber and contact) for production of H₂SO₄, and it is available commercially in a number of grades and concentrations. The lead chamber process, the older of the two processes, is used to produce much of the acid used to make fertilizers; it produces a relatively dilute acid (62%-78% H₂SO₄). The contact process produces a purer, more concentrated acid but requires purer raw materials and the use of expensive catalysts. In both processes sulfur dioxide is oxidized and dissolved in water. There are four steps to the contact process which starts with elemental sulfur.

1.      Step 1: The Burner. Pure sulfur is burned in air to give sulfur dioxide (SO₂), which is familiar to as rotten egg gas. This reaction happens at about 1000⁰C. The air used for the burning must be dust free as the presence of dust in the second part of the process will cause damage to the catalyst.

S_(g) + O_2(g)à SO_2(g)

2.      Step 2: The Converter. The sulfur dioxide is passed into a tower made of stacked vertical beds of catalyst material, usually vanadium oxide (V₂O₅) but sometimes platinum. Here the sulfur dioxide reacts with oxygen once again to give sulfur trioxide. The purpose of the catalyst is to reduce the temperature needed for the reaction to occur. This is used because the reaction of SO₂ with O₂is more productive at the lower temperature of 300⁰C.Sulfur Trioxide cannot be immediately converted into Sulfuric acid by absorbing water, as happens in acid rain. If allowed to absorb water the SO₃ becomes a fine mist that is very difficult to collect.

SO_2(g) + ½ O_2(g)à SO_3(g)

3.      Step 3: Absorption of sulfur trioxide. In order to produce sulfuric acid in liquid form, the sulfur trioxide gas from the converter is bubbled through very high purity liquid H₂SO₄. This is a very efficient process that results in virtually complete absorption of the SO₃ gas. The resulting liquid chemical is called oleum and has the formula H₂S₂O₇.

SO_3(g) + H₂SO_4(l)à H₂S₂O_7(l)

4.      Step 4: Hydration of the oleum. Finally the oleum is converted into highly concentrated H₂SO₄ by the addition of dilute H₂S₂O₇ and water. This reaction gives off a lot of heat and so the resulting acid needs to be cooled before it is sent to storage.

H₂S₂O_7(l) + H₂O_(l)à  H₂SO_4(aq)

The other well-known of the sulfur compound is hydrogen sulfide (H₂S). Hydrogen sulfide is colorless gas, and known as rotten egg smell gas. In chemistry terms, it can be said complete, if the structure and the chemical properties of hydrogen sulfide is explained. When the sulfur and the other chalcogenides make a bond with hydrogen, they form compounds with analogous formulas that may represented as H₂Y where Y, in this term is sulfur. The hydrogen chalconides molecules are angular, since there are four groups of electrons on each central chalcogens. As one of the chalcogen, the sulfur molecule is angular too. Snyder (1966) mention the properties of Hydrogen sulfide such as melting point, boiling point heat of fusion, heat of vaporization critical temperature dipole moment and dielectric constant. The melting point of hydrogen sulfide is 190 K and the boiling point of hydrogen sulfide is 213 K. Heat of fusion and heat of vaporization of hydrogen sulfide in a row is 0,568 kcal/mole and 4,46 kcal/mole. The dipole moment of hydrogen sulfide is 0,931 debyes, and the dielectric constant of hydrogen sulfide at 212 K is 8,0.

As said before, Hydrogen sulfide is colorless gas, and known as rotten egg odor gas. Hydrogen sulfide itself is a poisonus gas and its more poisonus than carbon monoxide, but its bad odor allows it to be detected at low concentrations. In nature, hydrogen sulfide is released into the air during volcanic eruption and by the decomposition of organic matter in the absence of air that’s rotten eggs smell of hydrogen sulfide. Hydrogen sulfide is difficult to dissolve in water and its solution become turbid because hydrogen sulfide is oxidized produce sulfur. In laboratory, hydrogen sulfide is used as strong reductor. The reaction is shown below.

H₂S + ½ O₂ à H₂O + SO₂

H₂S + SO₂ à 2 H₂O + 3S

Brady (1990), mention that hydrogen sulfide can be prepared in laboratory by reacting a metal sulfide with a strong monoxidizing acid, for example, HCl.

FeS_(s) + 2H⁺_(aq) à Fe²⁺_(aq) + H₂S_(g)

Another method, generally used when one wants to generate the H₂S in an aqueous solution, is the hydrolysis of an organic compound called thioacetamide.

Thioacetamide_(aq) + 2H₂O à H₂S_(aq) + NH₄⁺_(aq) + acetate ion_(aq)

The advantage of this reaction is that it avoids the release of significant amounts of toxic H₂S into the atmosphere. Qualitatively, hydrogen sulfide can be analyzed with paper which is tear-stained by salt solution of Pb²⁺ and produce black spot, because PbS is formed.

H₂S_(g) + Pb²⁺_(aq) à PbS + 2H⁺_(aq)

The last sulfur compound that will be explained is sodium thiosulfate (Na₂S₂O₃). Sodium thiosulfate is formed in reaction of sodium sulfites solutions and sulfur. The reaction is shown below.

S_(aq) + Na₂SO­_3(aq) à Na₂S₂O_3(aq)

The sodium thiosulfate has characteristic as reductor and used in iodometric analysis, according reaction:

2Na₂S₂O₃ + I₂ à Na₂S₄O₆ + 2 NaI

Specifically, the thiosulfates ion has the structure SSO₃^2-, and may be considered to be derived from sulfate by replacement of an atom by an S atom.

Based on the explanation above, it can be concluded that sulfur is common element that has several advantage and disadvantage for human being. Sulfur has allotropes. Sulfur exists in nature in pure element and in compound form. Some of sulfur compounds are sulfur dioxide, sulfuric acid, hydrogen sulfide and sodium thiosulfate. Every sulfur compound has their own properties, structure and process to prepare or making it. After knowing all about sulfur that is explained in this essay, the peoples understanding and its used is expected to be improve at the future.

REFERENCES

Brady, James E. 1990. General Chemistry Principles and Structure. Canada:Jhon Wiley and Sons, Inc.

Cotton, F. Albert., Wilkinson, Geoffrey., Gaus, Paul.L. 1995. Basic Inorganic Chemistry. Singapore: John Willey & Sons Pte. Ltd

Freshney, Paul Alan. 2007. PA Freshney Periodic Table v3.3.

Petrucci Hill, Ralph., Hill Jhon. 2002. General Chemistry An Integrated Approach. New Jersey:Prentice-Hall, Inc.

Snyder, Milton K. 1966. Chemistry Structure and Relations. USA: Holt, Rinehart and Winston, Inc.

Siregar, Manimpan., Sudria, Ida Bagus Nyoman. 2000. Buku Ajar Kimia Anorganik I. Singaraja : STKIP Negeri Singaraja