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Learning Objectives

• To understand the trends in properties and reactivity of the group 16 elements: the chalcogens.

The chalcogens are the first group in the p block to have no stable metallic elements. All isotopes of polonium (Po), the only metal in group 16, are radioactive, and only one element in the group, tellurium (Te), can even be described as a semimetal. As in groups 14 and 15, the lightest element of group 16, oxygen, is found in nature as the free element.

The atomic number of Sulfur (American) and Sulphur (British) is 8, it has 8 protons in the nucleus. Its most common isotope not found in the heart of supernovae also has 8 neutrons, making its atomic mass 16amu.Just to be clear, the atomic number is the number of protons contained in the nucleus. The reactions of sulfur atoms with propadiene and 1,2-butadiene. Canadian Journal of Chemistry (1985), 63(3), 667-75. Canadian Journal of Chemistry (1985), 63(3), 667-75. Human Metabolome Database (HMDB).

Of the group 16 elements, only sulfur was known in ancient times; the others were not discovered until the late 18th and 19th centuries. Sulfur is frequently found as yellow crystalline deposits of essentially pure S₈ in areas of intense volcanic activity or around hot springs. As early as the 15th century BC, sulfur was used as a fumigant in Homeric Greece because, when burned, it produces SO₂ fumes that are toxic to most organisms, including vermin hiding in the walls and under the floors of houses. Hence references to sulfur are common in ancient literature, frequently in the context of religious purification. In fact, the association of sulfur with the divine was so pervasive that the prefixes thio- (meaning “sulfur”) and theo- (meaning “god”) have the same root in ancient Greek. Though used primarily in the production of sulfuric acid, sulfur is also used to manufacture gunpowder and as a cross-linking agent for rubber, which enables rubber to hold its shape but retain its flexibility.

Atomic Number of Sulfur Atomic Number of Sulfur is 16. Chemical symbol for Sulfur is S. Number of protons in Sulfur is 16. Atomic weight of Sulfur is 32.06 u or g/mol. Melting point of Sulfur is 113 °C and its the boiling point is 444,7 °C.

Group 16 is the first group in the p block with no stable metallic elements.

Oxygen was not discovered until 1771, when the Swedish pharmacist Carl Wilhelm Scheele found that heating compounds such as KNO₃, Ag₂CO₃, and HgO produced a colorless, odorless gas that supported combustion better than air. The results were not published immediately, however, so Scheele’s work remained unknown until 1777. Unfortunately, this was nearly two years after a paper by the English chemist Joseph Priestley had been published, describing the isolation of the same gas by using a magnifying glass to focus the sun’s rays on a sample of HgO. Oxygen is used primarily in the steel industry during the conversion of crude iron to steel using the Bessemer process. Another important industrial use of oxygen is in the production of TiO₂, which is commonly used as a white pigment in paints, paper, and plastics.

Tellurium was discovered accidentally in 1782 by the Austrian chemist Franz Joseph Müller von Reichenstein, the chief surveyor of mines in Transylvania who was also responsible for the analysis of ore samples. The silvery-white metal had the same density as antimony but very different properties. Because it was difficult to analyze, Müller called it metallum problematicum (meaning “difficult metal”). The name tellurium (from the Latin tellus, meaning “earth”) was coined by another Austrian chemist, Martin Klaproth, who demonstrated in 1798 that Müller’s “difficult metal” was actually a new element. Tellurium is used to color glass and ceramics, in the manufacture of blasting caps, and in thermoelectric devices.

Jöns Jakob Berzelius (1779–1848)

Berzelius was born into a well-educated Swedish family, but both parents died when he was young. He studied medicine at the University of Uppsala, where his experiments with electroshock therapy caused his interests to turn to electrochemistry. Berzelius devised the system of chemical notation that we use today. In addition, he discovered six elements (cerium, thorium, selenium, silicon, titanium, and zirconium).

The heaviest chalcogen, polonium, was isolated after an extraordinary effort by Marie Curie. Although she was never able to obtain macroscopic quantities of the element, which she named for her native country of Poland, she demonstrated that its chemistry required it to be assigned to group 16. Marie Curie was awarded a second Nobel Prize in Chemistry in 1911 for the discovery of radium and polonium.

Preparation and General Properties of the Group 16 Elements

Oxygen is by far the most abundant element in Earth’s crust and in the hydrosphere (about 44% and 86% by mass, respectively). The same process that is used to obtain nitrogen from the atmosphere produces pure oxygen. Oxygen can also be obtained by the electrolysis of water, the decomposition of alkali metal or alkaline earth peroxides or superoxides, or the thermal decomposition of simple inorganic salts, such as potassium chlorate in the presence of a catalytic amount of MnO₂:

[mathrm{2KClO_3(s)overset{MnO_2(s)}{underset{Delta}rightleftharpoons}2KCl(s)+3O_2(g)} label{22.4.1}]

Unlike oxygen, sulfur is not very abundant, but it is found as elemental sulfur in rock formations overlying salt domes, which often accompany petroleum deposits (Figure (PageIndex{1})). Sulfur is also recovered from H₂S and organosulfur compounds in crude oil and coal and from metal sulfide ores such as pyrite (FeS₂).

Because selenium and tellurium are chemically similar to sulfur, they are usually found as minor contaminants in metal sulfide ores and are typically recovered as by-products. Even so, they are as abundant in Earth’s crust as silver, palladium, and gold. One of the best sources of selenium and tellurium is the “slime” deposited during the electrolytic purification of copper. Both of these elements are notorious for the vile odors of many of their compounds. For example, when the body absorbs even trace amounts of tellurium, dimethyltellurium [(CH₃)₂Te] is produced and slowly released in the breath and perspiration, resulting in an intense garlic-like smell that is commonly called “tellurium breath.”

With their ns²np⁴ electron configurations, the chalcogens are two electrons short of a filled valence shell. Thus in reactions with metals, they tend to acquire two additional electrons to form compounds in the −2 oxidation state. This tendency is greatest for oxygen, the chalcogen with the highest electronegativity. The heavier, less electronegative chalcogens can lose either four np electrons or four np and two ns electrons to form compounds in the +4 and +6 oxidation state, respectively, as shown in Table Figure (PageIndex{1}). As with the other groups, the lightest member in the group, in this case oxygen, differs greatly from the others in size, ionization energy, electronegativity, and electron affinity, so its chemistry is unique. Also as in the other groups, the second and third members (sulfur and selenium) have similar properties because of shielding effects. Only polonium is metallic, forming either the hydrated Po²⁺ or Po⁴⁺ ion in aqueous solution, depending on conditions.

Table (PageIndex{1}): Selected Properties of the Group 16 Elements

Property	Oxygen	Sulfur	Selenium	Tellurium	Polonium
*The configuration shown does not include filled d and f subshells.
†The values cited for the hexacations are for six-coordinate ions and are only estimated values.
atomic mass (amu)	16.00	32.07	78.96	127.60	209
atomic number	8	16	34	52	84
atomic radius (pm)	48	88	103	123	135
atomic symbol	O	S	Se	Te	Po
density (g/cm3) at 25°C	1.31 (g/L)	2.07	4.81	6.24	9.20
electron affinity (kJ/mol)	−141	−200	−195	−190	−180
electronegativity	3.4	2.6	2.6	2.1	2.0
first ionization energy (kJ/mol)	1314	1000	941	869	812
ionic radius (pm)†	140 (−2)	184 (−2), 29 (+6)	198 (−2), 42 (+6)	221 (−2), 56 (+6)	230 (−2), 97 (+4)
melting point/boiling point (°C)	−219/−183	115/445	221/685	450/988	254/962
normal oxidation state(s)	−2	+6, +4, −2	+6, +4, −2	+6, +4, −2	+2 (+4)
product of reaction with H2	H2O	H2S	H2Se	none	none
product of reaction with N2	NO, NO2	none	none	none	none
product of reaction with O2	—	SO2	SeO2	TeO2	PoO2
product of reaction with X2	O2F2	SF6, S2Cl2, S2Br2	SeF6, SeX4	TeF6, TeX4	PoF4, PoCl2, PoBr2
standard reduction potential (E°, V) (E0 → H2E in acidic solution)	+1.23	+0.14	−0.40	−0.79	−1.00
type of oxide	—	acidic	acidic	amphoteric	basic
valence electron configuration*	2s22p4	3s23p4	4s24p4	5s25p4	6s26p4

Reactions and Compounds of Oxygen

As in groups 14 and 15, the lightest group 16 member has the greatest tendency to form multiple bonds. Thus elemental oxygen is found in nature as a diatomic gas that contains a net double bond: O=O. As with nitrogen, electrostatic repulsion between lone pairs of electrons on adjacent atoms prevents oxygen from forming stable catenated compounds. In fact, except for O₂, all compounds that contain O–O bonds are potentially explosive. Ozone, peroxides, and superoxides are all potentially dangerous in pure form. Ozone (O₃), one of the most powerful oxidants known, is used to purify drinking water because it does not produce the characteristic taste associated with chlorinated water. Hydrogen peroxide (H₂O₂) is so thermodynamically unstable that it has a tendency to undergo explosive decomposition when impure:

[2H_2O_{2(l)} rightarrow 2H_2O_{(l)} + O_{2(g)} ;;; ΔG^o = −119; kJ/mol label{1}]

As in groups 14 and 15, the lightest element in group 16 has the greatest tendency to form multiple bonds.

Despite the strength of the O=O bond ((D_mathrm{O_2}) = 494 kJ/mol), (O_2) is extremely reactive, reacting directly with nearly all other elements except the noble gases. Some properties of O₂ and related species, such as the peroxide and superoxide ions, are in Table (PageIndex{2}). With few exceptions, the chemistry of oxygen is restricted to negative oxidation states because of its high electronegativity (χ = 3.4). Unlike the other chalcogens, oxygen does not form compounds in the +4 or +6 oxidation state. Oxygen is second only to fluorine in its ability to stabilize high oxidation states of metals in both ionic and covalent compounds. For example, AgO is a stable solid that contains silver in the unusual Ag(II) state, whereas OsO₄ is a volatile solid that contains Os(VIII). Because oxygen is so electronegative, the O–H bond is highly polar, creating a large bond dipole moment that makes hydrogen bonding much more important for compounds of oxygen than for similar compounds of the other chalcogens.

Table (PageIndex{2}): Some Properties of O₂ and Related Diatomic Species

Species	Bond Order	Number of Unpaired e−	O–O Distance (pm)*
*Source of data: Lauri Vaska, “Dioxygen-Metal Complexes: Toward a Unified View,” Accounts of Chemical Research 9 (1976): 175.
O2+	2.5	1	112
O2	2	2	121
O2−	1.5	1	133
O22−	1	0	149

Metal oxides are usually basic, and nonmetal oxides are acidic, whereas oxides of elements that lie on or near the diagonal band of semimetals are generally amphoteric. A few oxides, such as CO and PbO₂, are neutral and do not react with water, aqueous acid, or aqueous base. Nonmetal oxides are typically covalent compounds in which the bonds between oxygen and the nonmetal are polarized (E^δ+–O^δ−). Consequently, a lone pair of electrons on a water molecule can attack the partially positively charged E atom to eventually form an oxoacid. An example is reacting sulfur trioxide with water to form sulfuric acid:

[H_2O_{(l)} + SO_{3(g)} rightarrow H_2SO_{4(aq)} label{2}]

The oxides of the semimetals and of elements such as Al that lie near the metal/nonmetal dividing line are amphoteric, as we expect:

[Al_2O_{3(s)} + 6H^+_{(aq)} rightarrow 2Al^{3+}_{(aq)} + 3H_2O_{(l)} label{3}]

[Al_2O_{3(s)} + 2OH^−_{(aq)} + 3H_2O_{(l)} rightarrow 2Al(OH)^−_{4(aq)} label{4}]

Oxides of metals tend to be basic, oxides of nonmetals tend to be acidic, and oxides of elements in or near the diagonal band of semimetals are generally amphoteric.

Example (PageIndex{1})

For each reaction, explain why the given products form.

1. Ga₂O₃(s) + 2OH⁻(aq) + 3H₂O(l) → 2Ga(OH)₄⁻(aq)
2. 3H₂O₂(aq) + 2MnO₄⁻(aq) + 2H⁺(aq) → 3O₂(g) + 2MnO₂(s) + 4H₂O(l)
3. KNO₃(s) (xrightarrow{Delta}) KNO(s) + O₂(g)

Given: balanced chemical equations

Asked for: why the given products form

Strategy:

Classify the type of reaction. Using periodic trends in atomic properties, thermodynamics, and kinetics, explain why the observed reaction products form.

Solution:

The Atomic Number Of Sulfur

1. Gallium is a metal. We expect the oxides of metallic elements to be basic and therefore not to react with aqueous base. A close look at the periodic table, however, shows that gallium is close to the diagonal line of semimetals. Moreover, aluminum, the element immediately above gallium in group 13, is amphoteric. Consequently, we predict that gallium will behave like aluminum (Equation (ref{4})).
2. Hydrogen peroxide is an oxidant that can accept two electrons per molecule to give two molecules of water. With a strong oxidant, however, H₂O₂ can also act as a reductant, losing two electrons (and two protons) to produce O₂. Because the other reactant is permanganate, which is a potent oxidant, the only possible reaction is a redox reaction in which permanganate is the oxidant and hydrogen peroxide is the reductant. Recall that reducing permanganate often gives MnO₂, an insoluble brown solid. Reducing MnO₄⁻ to MnO₂ is a three-electron reduction, whereas the oxidation of H₂O₂ to O₂ is a two-electron oxidation.
3. This is a thermal decomposition reaction. Because KNO₃ contains nitrogen in its highest oxidation state (+5) and oxygen in its lowest oxidation state (−2), a redox reaction is likely. Oxidation of the oxygen in nitrate to atomic oxygen is a two-electron process per oxygen atom. Nitrogen is likely to accept two electrons because oxoanions of nitrogen are known only in the +5 (NO₃⁻) and +3 (NO₂⁻) oxidation states.

Exercise (PageIndex{2})

Predict the product(s) of each reaction and write a balanced chemical equation for each reaction.

1. SiO₂(s) + H⁺(aq) →
2. NO(g) + O₂(g) →
3. SO₃(g) + H₂O(l) →
4. H₂O₂(aq) + I^–(aq) →

Answer

1. SiO₂(s) + H⁺(aq) → no reaction
2. 2NO(g) + O₂(g) → 2NO₂(g)
3. SO₃(g) + H₂O(l) → H₂SO₄(aq)
4. H₂O₂(aq) + 2I⁻(aq) → I₂(aq) + 2OH⁻(aq)

Reactions and Compounds of the Heavier Chalcogens

Because most of the heavier chalcogens (group 16) and pnicogens (group 15) are nonmetals, they often form similar compounds. For example, both third-period elements of these groups (phosphorus and sulfur) form catenated compounds and form multiple allotropes. Consistent with periodic trends, the tendency to catenate decreases as we go down the column.

Sulfur and selenium both form a fairly extensive series of catenated species. For example, elemental sulfur forms S₈ rings packed together in a complex “crankshaft” arrangement (Figure (PageIndex{2})), and molten sulfur contains long chains of sulfur atoms connected by S–S bonds. Moreover, both sulfur and selenium form polysulfides (S_n²⁻) and polyselenides (Se_n²⁻), with n ≤ 6. The only stable allotrope of tellurium is a silvery white substance whose properties and structure are similar to those of one of the selenium allotropes. Polonium, in contrast, shows no tendency to form catenated compounds. The striking decrease in structural complexity from sulfur to polonium is consistent with the decrease in the strength of single bonds and the increase in metallic character as we go down the group.

As in group 15, the reactivity of elements in group 16 decreases from lightest to heaviest. For example, selenium and tellurium react with most elements but not as readily as sulfur does. As expected for nonmetals, sulfur, selenium, and tellurium do not react with water, aqueous acid, or aqueous base, but all dissolve in strongly oxidizing acids such as HNO₃ to form oxoacids such as H₂SO₄. In contrast to the other chalcogens, polonium behaves like a metal, dissolving in dilute HCl to form solutions that contain the Po²⁺ ion.

Just as with the other groups, the tendency to catenate, the strength of single bonds, and reactivity decrease down the group.

Fluorine reacts directly with all chalcogens except oxygen to produce the hexafluorides (YF₆), which are extraordinarily stable and unreactive compounds. Four additional stable fluorides of sulfur are known; thus sulfur oxidation states range from +1 to +6 (Figure (PageIndex{2})). In contrast, only four fluorides of selenium (SeF₆, SeF₄, FSeSeF, and SeSeF₂) and only three of tellurium (TeF₄, TeF₆, and Te₂F₁₀) are known.

Direct reaction of the heavier chalcogens with oxygen at elevated temperatures gives the dioxides (YO₂), which exhibit a dramatic range of structures and properties. The dioxides become increasingly metallic in character down the group, as expected, and the coordination number of the chalcogen steadily increases. Thus SO₂ is a gas that contains V-shaped molecules (as predicted by the valence-shell electron-pair repulsion model), SeO₂ is a white solid with an infinite chain structure (each Se is three coordinate), TeO₂ is a light yellow solid with a network structure in which each Te atom is four coordinate, and PoO₂ is a yellow ionic solid in which each Po⁴⁺ ion is eight coordinate.

The dioxides of sulfur, selenium, and tellurium react with water to produce the weak, diprotic oxoacids (H₂YO₃—sulfurous, selenous, and tellurous acid, respectively). Both sulfuric acid and selenic acid (H₂SeO₄) are strong acids, but telluric acid [Te(OH)₆] is quite different. Because tellurium is larger than either sulfur or selenium, it forms weaker π bonds to oxygen. As a result, the most stable structure for telluric acid is Te(OH)₆, with six Te–OH bonds rather than Te=O bonds. Telluric acid therefore behaves like a weak triprotic acid in aqueous solution, successively losing the hydrogen atoms bound to three of the oxygen atoms. As expected for compounds that contain elements in their highest accessible oxidation state (+6 in this case), sulfuric, selenic, and telluric acids are oxidants. Because the stability of the highest oxidation state decreases with increasing atomic number, telluric acid is a stronger oxidant than sulfuric acid.

The stability of the highest oxidation state of the chalcogens decreases down the column.

Sulfur and, to a lesser extent, selenium react with carbon to form an extensive series of compounds that are structurally similar to their oxygen analogues. For example, CS₂ and CSe₂ are both volatile liquids that contain C=S or C=Se bonds and have the same linear structure as CO₂. Because these double bonds are significantly weaker than the C=O bond, however, CS₂, CSe₂, and related compounds are less stable and more reactive than their oxygen analogues. The chalcogens also react directly with nearly all metals to form compounds with a wide range of stoichiometries and a variety of structures. Metal chalcogenides can contain either the simple chalcogenide ion (Y²⁻), as in Na₂S and FeS, or polychalcogenide ions (Y_n²⁻), as in FeS₂ and Na₂S₅.

The dioxides of the group 16 elements become increasingly basic, and the coordination number of the chalcogen steadily increases down the group.

Ionic chalcogenides like Na₂S react with aqueous acid to produce binary hydrides such as hydrogen sulfide (H₂S). Because the strength of the Y–H bond decreases with increasing atomic radius, the stability of the binary hydrides decreases rapidly down the group. It is perhaps surprising that hydrogen sulfide, with its familiar rotten-egg smell, is much more toxic than hydrogen cyanide (HCN), the gas used to execute prisoners in the “gas chamber.” Hydrogen sulfide at relatively low concentrations deadens the olfactory receptors in the nose, which allows it to reach toxic levels without detection and makes it especially dangerous.

Example (PageIndex{2})

For each reaction, explain why the given product forms or no reaction occurs.

1. SO₂(g) + Cl₂(g) → SO₂Cl₂(l)
2. SF₆(g) + H₂O(l) → no reaction
3. 2Se(s) + Cl₂(g) → Se₂Cl₂(l)

Given: balanced chemical equations

Asked for: why the given products (or no products) form

Strategy:

Classify the type of reaction. Using periodic trends in atomic properties, thermodynamics, and kinetics, explain why the reaction products form or why no reaction occurs.

Solution:

1. One of the reactants (Cl₂) is an oxidant. If the other reactant can be oxidized, then a redox reaction is likely. Sulfur dioxide contains sulfur in the +4 oxidation state, which is 2 less than its maximum oxidation state. Sulfur dioxide is also known to be a mild reducing agent in aqueous solution, producing sulfuric acid as the oxidation product. Hence a redox reaction is probable. The simplest reaction is the formation of SO₂Cl₂ (sulfuryl chloride), which is a tetrahedral species with two S–Cl and two S=O bonds.

2. Sulfur hexafluoride is a nonmetallic halide. Such compounds normally react vigorously with water to produce an oxoacid of the nonmetal and the corresponding hydrohalic acid. In this case, however, we have a highly stable species, presumably because all of sulfur’s available orbitals are bonding orbitals. Thus SF₆ is not likely to react with water.
3. Here we have the reaction of a chalcogen with a halogen. The halogen is a good oxidant, so we can anticipate that a redox reaction will occur. Only fluorine is capable of oxidizing the chalcogens to a +6 oxidation state, so we must decide between SeCl₄ and Se₂Cl₂ as the product. The stoichiometry of the reaction determines which of the two is obtained: SeCl₄ or Se₂Cl₂.

Exercise (PageIndex{2})

Predict the products of each reaction and write a balanced chemical equation for each reaction.

1. Te(s) + Na(s) (xrightarrow{Delta})
2. SF₄(g) + H₂O(l) →
3. CH₃SeSeCH₃(soln) + K(s) →
4. Li₂Se(s) + H⁺(aq) →

Answer

1. Te(s) + 2Na(s) → Na₂Te(s)
2. SF₄(g) + 3H₂O(l) → H₂SO₃(aq) + 4HF(aq)
3. CH₃SeSeCH₃(soln) + 2K(s) → 2KCH₃Se(soln)
4. Li₂Se(s) + 2H⁺(aq) → H₂Se(g) + 2Li⁺(aq)

Summary

The chalcogens have no stable metallic elements. The tendency to catenate, the strength of single bonds, and the reactivity all decrease moving down the group. Because the electronegativity of the chalcogens decreases down the group, so does their tendency to acquire two electrons to form compounds in the −2 oxidation state. The lightest member, oxygen, has the greatest tendency to form multiple bonds with other elements. It does not form stable catenated compounds, however, due to repulsions between lone pairs of electrons on adjacent atoms. Because of its high electronegativity, the chemistry of oxygen is generally restricted to compounds in which it has a negative oxidation state, and its bonds to other elements tend to be highly polar. Metal oxides are usually basic, and nonmetal oxides are acidic, whereas oxides of elements along the dividing line between metals and nonmetals are amphoteric. The reactivity, the strength of multiple bonds to oxygen, and the tendency to form catenated compounds all decrease down the group, whereas the maximum coordination numbers increase. Because Te=O bonds are comparatively weak, the most stable oxoacid of tellurium contains six Te–OH bonds. The stability of the highest oxidation state (+6) decreases down the group. Double bonds between S or Se and second-row atoms are weaker than the analogous C=O bonds because of reduced orbital overlap. The stability of the binary hydrides decreases down the group.

Home » Minerals » Sulfur

Chemical element. Native mineral. Essential to all living things.

Article by: Hobart M. King, PhD, RPG

Sulfur terminal: Piles of yellow sulfur at a terminal near Vancouver, British Columbia, Canada. The sulfur is brought by rail from oil and natural gas processing facilities in the Province of Alberta. At this terminal it is loaded onto barges and ships for bulk transport. Photo copyright iStockphoto / teekaygee.

Sulfur fumarole: As hot volcanic gases, rich in sulfur, escape from a volcanic vent, the gases cool and sulfur is deposited as yellow crystals around the vent. This fumarole on the island of Kunashir (in the Kuril Islands, northeast of the Japanese island of Hokkaido) has a significant accumulation of bright yellow sulfur. Photo copyright iStockphoto / Sergey Dubrovskiy.

Did You Know? Many strong odors are produced by sulfur compounds. The smell of skunks, matches, garlic, grapefruit, and rotten eggs are caused by sulfur. Image copyright iStockphoto / Florintt, Gio_banfi, Abomb Industries Design, ivelly, and Big_Ryan.

What is Sulfur?

Sulfur is a chemical element with an atomic number of 16 and an atomic symbol of S. At room temperature it is a yellow crystalline solid. Even though it is insoluble in water, it is one of the most versatile elements at forming compounds. Sulfur reacts and forms compounds with all elements except gold, iodine, iridium, nitrogen, platinum, tellurium, and the inert gases.

Sulfur is abundant and occurs throughout the Universe, but it is rarely found in a pure, uncombined form at Earth's surface. As an element, sulfur is an important constituent of sulfate and sulfide minerals. It occurs in the dissolved ions of many waters. It is an important constituent of many atmospheric, subsurface, and dissolved gases. It is an essential element in all living things and is in the organic molecules of all fossil fuels.

Did You Know? The Chinese discovered sulfur in about 2000 BC, used it to make gunpowder in the 7th century, and used gunpowder to launch rockets, shoot projectiles, and make hand grenades in the 10th century.

Physical Properties of Sulfur
Chemical Classification	Native element
Color	Yellow. Brownish yellow to greenish yellow. Red when molten at over 200 degrees Celsius. Burns with a flame that can be difficult to see in daylight but is blue in the dark.
Streak	Yellow
Luster	Crystals are resinous to greasy. Powdered sulfur is dull or earthy.
Diaphaneity	Transparent to translucent
Cleavage	None
Mohs Hardness	1.5 to 2.5
Specific Gravity	2.0 to 2.1
Diagnostic Properties	Yellow color, low hardness, low specific gravity, extremely flammable burning with a blue flame, low melting temperature
Chemical Composition	S
Crystal System	Orthorhombic
Uses	About 90% is used to manufacture sulfuric acid. The remainder is used in a variety of products that include hydrogen sulfide, insecticides, herbicides, fungicides, pharmaceuticals, soaps, textiles, papers, processed rubber, gunpowder, leather, paint, dyes, food preservatives.

World Sulfur Production: During 2015, an estimated 70 million metric tons of sulfur was produced worldwide. The production was widely divided among a large number of countries. The top 12 producing countries were China, the United States, Russia, Canada, Germany, Japan, Saudi Arabia, India, Kazakhstan, Iran, United Arab Emirates, and Mexico. These countries are where the sulfur was separated from its geologic source material rather than the original source of the sulfur, since most sulfur is separated when fossil fuels are processed or sulfide ores are smelted. Data from the United States Geological Survey. [7]

Sulfur is Abundant and Everywhere!

The information below should convince you that sulfur is extremely abundant and present everywhere.

• 11th most abundant element in the human body [1]
• 6th most abundant element in seawater [2]
• 14th most abundant element in Earth’s crust [3]
• 9th most abundant element in the entire Earth [4]
• 10th most abundant element in the solar system [5]
• 10th most abundant element in the Universe [6]

Sulfur Crystals: Bright yellow sulfur crystal group showing the mineral's characteristic orthorhombic crystal form and resinous luster. Specimen measures approximately 7.3 x 6.6 x 5.3 centimeters in size and was collected from the Agrigento Province, Sicily, Italy. Specimen and photo by Arkenstone / www.iRocks.com.

Burning sulfur: Pieces of sulfur burning in daylight and in the dark. Photo by Johannes 'volty' Hemmerlein, used here under a GNU Free Documentation License.

Did You Know? Jupiter's moon, Io, has over 400 active volcanoes that emit enormous amounts of sulfur - so much sulfur that the moon has a yellowish color.

'Sulfur' or 'Sulphur'?

The name 'sulphur' has been used in the United Kingdom and throughout the British Empire for hundreds of years. 'Sulfur' is the spelling used in common and scientific communication in the United States. In 1990 the International Union of Pure and Applied Chemistry designated 'sulfur' as the preferred spelling. How the word is spelled can often reveal the age and origin of publications and authors.

Information Sources

[1] What Elements Are Found in the Human Body? Article in the Building Blocks of Life section of the Arizona School of Life Sciences website, accessed November 2016. [2] Periodic Table of Elements in the Ocean, article on the Monterey Bay Aquarium Research Institute website, accessed November 2016. [3] List of Periodic Table Elements Sorted by Abundance in Earth's Crust, article on the Israel Science and Technology website, accessed November 2016. [4] The Composition of the Earth, by William F. McDonough, Chapter 1 in Earthquake Thermodynamics and Phase Transformations in the Earth’s Interior, manuscript on the Massachusetts Institute of Technology website, accessed November 2016. [5] Solar System Abundances and Condensation Temperatures of the Elements by Katharina Lodders, article published on The Astrophysical Journal website, accessed November 2016. [6] Abundance in the Universe of the Elements, article on the PeriodicTable.com website, accessed November 2016. [7] Sulfur, by Lori E. Apodaca, United States Geological Survey, Mineral Commodity Summaries, 2016. [8] The International Mineralogical Association Database of Mineral Properties, an online database of minerals along with their chemical and physical properties that can be queried and sorted by anyone with internet access.

Sulfur as a Native Element Mineral

As a mineral, sulfur is a bright yellow crystalline material. It forms near volcanic vents and fumaroles, where it sublimates from a stream of hot gases. Small amounts of native sulfur also form during the weathering of sulfate and sulfide minerals.

The largest accumulations of mineral sulfur are found in the subsurface. Many of these are in fractures and cavities associated with sulfide ore mineralization. The largest are associated with evaporite minerals, where gypsum and anhydrite yield native sulfur as a product of bacterial action. Significant amounts of sulfur have been produced from the cap rock of salt domes but this type of production is rarely done today.

Related: Blue Flames caused by burning sulfur, in this night scene from Kawah Ijen Volcano, located on the island of Java, Indonesia.

Atomic Number Of Sulfur Atom

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Minerals That Contain Sulfur

According to the International Mineralogical Association's database, over 1000 minerals contain sulfur as an essential part of their composition. [8] This is a result of sulfur's ability to form compounds with all but a few other elements. The tables below list a small number of sulfide, sulfarsenide, sulfosalt and sulfate minerals. Many of the most common sulfur minerals are included in the list, but the list is not intended to be complete.

Sulfide Minerals:
Mineral	Composition
Acanthite	Ag2S
Chalcocite	Cu2S
Bornite	Cu5FeS4
Galena	PbS
Sphalerite	ZnS
Chalcopyrite	CuFeS2
Pyrrhotite	Fe1-xS
Millerite	NiS
Pentlandite	(Fe,Ni)9S8
Covellite	CuS
Cinnabar	HgS
Realgar	AsS
Orpiment	As2S3
Stibnite	Sb2S3
Pyrite	FeS2
Marcasite	FeS2
Molybdenite	MoS2

What Is The Atomic Number Of Sulfur 4

Sulfarsenide Minerals:
Mineral	Composition
Cobaltite	(Co,Fe)AsS
Arsenopyrite	FeAsS
Gersdorffite	NiAsS

Sulfosalt Minerals:
Mineral	Composition
Pyrargyrite	Ag3SbS3
Proustite	Ag3AsS3
Tetrahedrite	Cu12Sb4S13
Tennantite	Cu12As4S13
Enargite	Cu3AsS4
Bournonite	PbCuSbS3
Jamesonite	Pb4FeSb6S14
Cylindrite	Pb3Sn4FeSb2S14

Atomic Number Of Sulfur

Hydroxide and Hydrous Sulfate Minerals:
Mineral	Composition
Gypsum	CaSO4·2H2O
Chalcanthite	CuSO4·5H2O
Kieserite	MgSO4·H2O
Starkeyite	MgSO4·4H2O
Hexahydrite	MgSO4·6H2O
Epsomite	MgSO4·7H2O
Meridianiite	MgSO4·11H2O
Melanterite	FeSO4·7H2O
Antlerite	Cu3SO4(OH)4
Brochantite	Cu4SO4(OH)6
Alunite	KAl3(SO4)2(OH)6
Jarosite	KFe3(SO4)2(OH)6

Anhydrous Sulfate Minerals:
Mineral	Composition
Barite	BaSO4
Celestite	SrSO4
Anglesite	PbSO4
Anhydrite	CaSO4
Hanksite	Na22K(SO4)9(CO3)2Cl

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Sulfur Atomic Model

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