18.10 Occurrence, Preparation, and Properties of Sulfur

Occurrence and Extraction of Sulfur

Sulfur is one of the few elements found in pure, elemental form in nature. It's also a major industrial chemical, with most of the world's supply going toward sulfuric acid production. Understanding where sulfur comes from and how it's extracted connects directly to both geology and industrial chemistry.

Natural Occurrence

Sulfur shows up in three main settings:

• Elemental sulfur deposits near volcanoes and underground salt domes (classic examples include Sicilian volcanoes and Gulf Coast salt domes in the U.S.)
• Sulfide minerals, where sulfur is bonded to metals: pyrite ($FeS_2$), galena ($PbS$), and sphalerite ($ZnS$)
• Fossil fuels, where sulfur is present mainly as hydrogen sulfide ($H_2S$) dissolved in natural gas and petroleum

Extraction Methods

The Frasch Process is used to pull elemental sulfur out of deep underground deposits. It works because sulfur has a relatively low melting point ($115°C$):

1. Superheated water (around $160°C$) is pumped down into the sulfur deposit, melting the sulfur in place.
2. Compressed air is then forced into the well, pushing the molten sulfur upward to the surface.
3. The molten sulfur is collected and allowed to cool and solidify. The product is quite pure (often 99.5%+).

Desulfurization and the Claus Process account for a large share of sulfur production today. When $H_2S$ is removed from natural gas or petroleum during refining, it gets converted to elemental sulfur through the Claus process. As environmental regulations have tightened, this recovered sulfur has actually become the dominant global source.

Allotropes and Reactivity of Sulfur

Sulfur is famous for its polymorphism, meaning it exists in multiple solid crystal forms (allotropes). The most important ones to know:

• Rhombic sulfur ($\alpha$-sulfur): The most common and thermodynamically stable form below $95.5°C$. It has an orthorhombic crystal structure, and its molecules are $S_8$ rings.
• Monoclinic sulfur ($\beta$-sulfur): Stable between $95.5°C$ and $119°C$. It forms needle-like crystals, also built from $S_8$ rings but packed differently.
• Amorphous sulfur: Formed when molten sulfur is cooled rapidly (like pouring it into cold water). The result is a rubbery, elastic solid. This form is used in rubber vulcanization.

The temperature $95.5°C$ is the transition temperature between rhombic and monoclinic forms. Below it, monoclinic sulfur slowly converts back to rhombic.

Chemical Reactivity

Elemental sulfur is moderately reactive and participates in several important reaction types:

• Combustion: Burns in air to produce sulfur dioxide.

$S_{(s)} + O_{2(g)} \rightarrow SO_{2(g)}$

• Reaction with metals: Combines directly with many metals to form sulfides.

$Fe_{(s)} + S_{(s)} \rightarrow FeS_{(s)}$

• Oxidation by strong oxidizing agents: For example, hydrogen peroxide can oxidize sulfur all the way to sulfate.

S_{(s)} + 4H_2O_{2(aq)} \rightarrow SO_4^{2-}_{(aq)} + 4H_2O_{(l)}

Sulfur vs Other Group 16 Elements

Sulfur sits in Group 16 alongside oxygen, selenium, and tellurium. Comparing their behavior highlights periodic trends.

• Sulfur exhibits a wide range of oxidation states: -2 (sulfide, as in $H_2S$), 0 (elemental), +4 (sulfite, $SO_3^{2-}$), and +6 (sulfate, $SO_4^{2-}$).
• Oxygen is more limited, typically showing only -2 (oxide) and 0 (elemental). Its small atomic size and lack of accessible d orbitals prevent it from reaching +4 or +6 states.
• Selenium and tellurium share the same set of oxidation states as sulfur (-2, 0, +4, +6). As you move down the group, the larger atomic size actually makes the +6 state more stable. Tellurium holds onto a +6 oxidation state more readily than sulfur does.

Chemical Properties and Bonding

Sulfur's chemistry revolves around its ability to both gain and lose electrons depending on what it's reacting with:

• In reduction reactions, sulfur gains electrons and drops to a -2 oxidation state (forming sulfides like $H_2S$).
• In oxidation reactions, sulfur loses electrons and can reach +4 or +6 states (forming compounds like $SO_2$ or $H_2SO_4$).

A distinctive feature of sulfur is catenation, the ability of sulfur atoms to bond to each other in long chains and rings. This is why sulfur forms $S_8$ rings in its solid allotropes and why polysulfide chains ($S_n^{2-}$) exist in solution. Catenation is possible because the S–S single bond is reasonably strong (about $266 \text{ kJ/mol}$), giving sulfur a richer structural chemistry than most nonmetals.