Part A: The Essentials
Group 16 elements are all nonmetals, apart from polonium, the heaviest member of the group. The group contains two of the most important elements for life. Oxygen is essential for higher life forms, both in the atmosphere and as water. Sulfur is essential to all life forms, and even selenium is required in trace amounts.
The Group 16 elements oxygen, sulfur, selenium, tellurium, and polonium are often called the chalcogens. The name derives from the Greek word for 'bronze', referring to the association of sulfur and its congeners with copper in metal ores.
Electronic Configuration
The group electron configuration of ns²np⁴ suggests a group maximum oxidation number of +6. Oxygen never achieves this maximum oxidation state, although the other elements do in some circumstances. The electron configuration also suggests that stability may be achieved with an oxidation number of −2, which is overwhelmingly common for O.
The most remarkable feature of S, Se, and Te is that they form stable compounds with oxidation numbers between −2 and +6.
16.1 The Elements
Oxygen, sulfur, and selenium are nonmetals, tellurium is a metalloid, and polonium is a metal. Allotropy and polymorphism are important features of Group 16, and sulfur occurs in more naturally occurring allotropes and polymorphs than any other element.
Selected Properties of the Elements
| Property | O | S | Se | Te | Po |
|---|---|---|---|---|---|
| Covalent radius /pm | 74 | 104 | 117 | 137 | 140 |
| Ionic radius /pm | 140 | 184 | 198 | 221 | — |
| First ionization energy /kJ mol⁻¹ | 1310 | 1000 | 941 | 870 | 812 |
| Melting point /°C | −218 | 113 (α) | 217 | 450 | 254 |
| Boiling point /°C | −183 | 445 | 685 | 990 | 960 |
| Pauling electronegativity | 3.4 | 2.6 | 2.6 | 2.1 | 2.0 |
| Electron affinity /kJ mol⁻¹ | 141 | 200 | 195 | 190 | 183 |
Oxygen - Abundance and Properties
Oxygen is the most abundant element in the Earth's crust at 46 per cent by mass, and is present in all silicate minerals. It comprises 86 per cent by mass of the oceans and 89 per cent of water. The average human is two-thirds oxygen by mass.
Allotropes of Oxygen
Sulfur Allotropes
Sulfur can exist in a large number of allotropic forms due to the high S−S bond energy of 265 kJ mol⁻¹, which is exceeded only by C−C (330 kJ mol⁻¹) and H−H (436 kJ mol⁻¹). All the crystalline forms of sulfur that can be isolated at room temperature consist of Sn rings.
S₈ Crown-like Ring Structure
| Allotrope | Melting Point /°C | Appearance |
|---|---|---|
| S₃ | Gas | Cherry red |
| S₆ | 50 (decomp.) | Orange-red |
| α-S₈ | 113 | Yellow |
| β-S₈ | 119 | Yellow |
| γ-S₈ | 107 | Pale yellow |
| S₁₂ | 148 | Pale yellow |
| S∞ | 104 | Yellow |
16.2 Simple Compounds
Water - The Most Important Hydride
Water is the only Group 16 hydride that is not a poisonous, malodorous gas. Its melting and boiling points (0°C and 100°C) are both very high compared to compounds of similar molecular mass due to extensive hydrogen bonding between H and the highly electronegative oxygen, O−H⋯O.
Properties of Group 16 Hydrides
| Property | H₂O | H₂S | H₂Se | H₂Te | H₂Po |
|---|---|---|---|---|---|
| Melting point /°C | 0.0 | −85.6 | −65.7 | −51 | −36 |
| Boiling point /°C | 100.0 | −60.3 | −41.3 | −4 | 37 |
| ΔfH° /kJ mol⁻¹ | −285.6 (l) | −20.1 | +73.0 | +99.6 | — |
| Bond length /pm | 96 | 134 | 146 | 169 | — |
| Bond angle /° | 104.5 | 92.1 | 91 | 90 | — |
| pKa1 | 14.00 | 6.89 | 3.89 | 2.64 | — |
Oxides and Sulfides
Oxygen forms oxides with most metals, and peroxides and superoxides with Group 1 and 2 metals. Sulfur forms sulfides (S²⁻) and disulfides (S₂²⁻) with metals. Selenium and tellurium form selenides and tellurides containing Se²⁻ and Te²⁻.
Sulfur Oxides
Sulfur dioxide is manufactured on a huge scale for the production of sulfuric acid via the contact process, where it is first oxidized to SO₃.
16.3 Ring and Cluster Compounds
Polythionic Acids
Sulfur forms many polythionic acids, H₂SnO₆, with up to six S atoms:
- Tetrathionate S₄O₆²⁻
- Pentathionate S₅O₆²⁻
Polysulfides
Many polysulfides of electropositive elements contain Sn²⁻ ions where n = 2−6.
Square-Planar E₄²⁺ Ions
The particular stability of the square-planar ions E₄²⁺ (E = S, Se, Te) is explained by molecular orbital theory. Each E atom has six valence electrons, giving 24 − 2 = 22 electrons in all. There are two lone pairs on each E atom, leaving six electrons to occupy molecular orbitals: one bonding, two nonbonding, and one antibonding (vacant).
Sulfur-Nitrogen Compounds
16.4 Oxygen
Dioxygen Production
Dioxygen is a biogenic gas — almost all of it is the result of photosynthesis. Oxygen is obtained on a massive scale by the liquefaction and distillation of liquid air for steelmaking.
Chloride process for TiO₂ production
Properties of Liquid Oxygen
Liquid oxygen is very pale blue and boils at −183°C. Its colour arises from electronic transitions involving pairs of neighbouring molecules: one photon can raise two O₂ molecules to an excited state.
Molecular Orbital Description
The MO description of O₂ implies a double bond. However, the outermost two electrons occupy different antibonding π orbitals with parallel spins, making the molecule paramagnetic.
Ground state: ³Σg⁻ (triplet, paramagnetic)
Singlet Oxygen
The singlet state ¹Δg, with both electrons paired in one π* orbital, lies 94 kJ mol⁻¹ above the ground state. O₂(¹Δg) survives long enough to participate in chemical reactions and reacts as an electrophile.
Ozone
Ozone (O₃) is an explosive and highly reactive endoergic blue gas (ΔfG° = +163 kJ mol⁻¹). The O₃ molecule is angular with a bond angle of 117° and is diamagnetic.
Ozone is exceeded in oxidizing power only by F₂, atomic O, the OH radical, and perxenate ions.
16.5 Reactivity of Oxygen
Oxygen is a strong oxidant, yet most of its reactions are sluggish. For example, a solution of Fe²⁺ is only slowly oxidized by air, even though the reaction is thermodynamically favourable.
Factors Contributing to Slow Kinetics
- High O=O bond energy (494 kJ mol⁻¹) — results in high activation energy for homolytic dissociation
- Triplet ground state — O₂ is neither an effective Lewis acid nor Lewis base
- Spin restriction — reactions of O₂ require appropriate spin states
Oxygen is an essential by-product of hydrogen generation by electrochemical water splitting. The direct oxidation of water to O₂ (E° = 1.23 V) is kinetically challenging because it requires removal of four protons and four electrons from two water molecules.
In biology, photosynthetic O₂ evolution occurs at a Mn−O cluster that can produce more than 100 molecules of O₂ per second. Promising non-biological electrocatalysts are based on cobalt oxides.
16.6 Sulfur
Extraction Methods
Underground deposits are forced to the surface using superheated water and steam, and compressed air. The process is energy-intensive and depends on access to cheap water and energy.
Extraction from natural gas and crude oil. H₂S is first oxidized in air at 1000−1400°C, producing some SO₂ that then reacts with remaining H₂S at 200−350°C over a catalyst:
Catenation in Sulfur
Unlike O, sulfur tends to form single bonds with itself rather than double bonds. This tendency arises from the relative strengths of p−p σ bonding (increases from O to S) and p−p π bonding (decreases). As a result, S aggregates into larger molecules and is a solid at room temperature.
Polymorphs of Sulfur
- α-S₈ (Orthorhombic) — common yellow form, crown-like eight-membered rings
- β-S₈ (Monoclinic) — forms when heated to 93°C
- γ-S₈ (Monoclinic) — forms when molten sulfur heated above 150°C is cooled slowly
Reactions of Sulfur
Industrial Uses
Most sulfur produced industrially is used to manufacture sulfuric acid, one of the most important manufactured chemicals. Other uses include:
- Synthesis of fertilizers
- Electrolyte in lead-acid batteries (dilute H₂SO₄)
- Component of gunpowder (KNO₃ + C + S)
- Vulcanization of natural rubber
16.7 Selenium, Tellurium, and Polonium
Selenium
Selenium can be extracted from waste sludge produced by sulfuric acid plants or from copper sulfide ores. It exists in several forms:
- Red selenium — contains Se₈ rings (α, β, γ forms)
- Metallic grey selenium — most stable form, composed of helical chains
- Amorphous black selenium — common commercial form with rings up to 1000 Se atoms
Selenium exhibits both photovoltaic (light → electricity) and photoconductive character due to its small band gap (2.6 eV crystalline, 1.8 eV amorphous).
Tellurium
Tellurium crystallizes in a chain structure like grey selenium.
Polonium
Polonium crystallizes in a primitive cubic structure — the only element to adopt this structure under normal conditions. It is highly toxic due to intense radioactivity — mass for mass, it is about 2.5 × 10¹¹ times as toxic as hydrocyanic acid.
16.8 Hydrides
(a) Water
At least nine distinct forms of ice have been identified. Water is formed by direct interaction of the elements:
(b) Hydrogen Peroxide
H₂O₂ is a very pale blue, viscous liquid with b.p. 150°C and density 1.445 g cm⁻³ at 25°C. It is unstable with respect to disproportionation:
H₂O₂ is a very powerful oxidizing agent in acid solution:
Fenton Reaction
The hydroxyl radical is one of the strongest oxidizing agents known (E = +2.85 V).
Question: Is Pd²⁺ thermodynamically capable of catalysing the decomposition of H₂O₂?
Answer: For the first half-reaction (Pd²⁺ → Pd + O₂), E°cell = +0.22 V. For the second half-reaction (Pd + H₂O₂ → Pd²⁺ + H₂O), E°cell = +0.84 V. Both are spontaneous (K > 1), so catalytic decomposition is thermodynamically favoured.
(c) Hydrogen Sulfide
H₂S is toxic, made more hazardous by the fact that it anaesthetizes the olfactory nerves. It can be prepared by:
H₂S is a weak acid:
16.9 Halides
Sulfur Halides
| Oxidation Number | Formula | Remarks |
|---|---|---|
| +½ | Te₂X (X = Br, I) | Silver-grey, halide bridges |
| +1 | S₂F₂, S₂Cl₂ | Two isomers, reactive |
| +2 | SCl₂ | Reactive |
| +4 | SF₄, SeX₄, TeX₄ | SF₄ gas, SeF₄ liquid, TeF₄ solid |
| +5 | S₂F₁₀, Se₂F₁₀ | Reactive |
| +6 | SF₆, SeF₆, TeF₆ | SF₆, SeF₆ colourless gases; TeF₆ liquid |
Sulfur Hexafluoride
SF₆ is a gas at room temperature and is very unreactive. Its inertness stems from steric protection of the central S atom. In contrast, SeF₆ is easily hydrolysed and more reactive.
(thermodynamically favourable but kinetically suppressed)
Sulfur Chlorides
Disulfur dichloride (S₂Cl₂) is a foul-smelling, toxic yellow liquid (b.p. 138°C). S₂Cl₂ and sulfur dichloride (SCl₂) are used in the vulcanization of rubber, where S-atom bridges are introduced between polymer chains.
16.10 Metal Oxides
The O₂ molecule readily removes electrons from metals to form oxides containing O²⁻ (oxide), O₂⁻ (superoxide), and O₂²⁻ (peroxide).
Structural Trends
- M(I): M₂O oxides — often rutile or antifluorite structure (6:3 and 8:4 coordination)
- M(II): MO oxides — usually rock-salt structure (6:6 coordination)
- M(III): M₂O₃ oxides — often 6:4 coordination
- MO₄ compounds — molecular (e.g., tetrahedral OsO₄)
16.11 Metal Sulfides, Selenides, Tellurides, and Polonides
Many metals occur naturally as their sulfide ores. The sulfides can be prepared by:
Layered Disulfides
Layered disulfides are built from a sulfide layer, a metal layer, and then another sulfide layer. These sandwiches stack together with sulfide layers adjacent. This is not consistent with a simple ionic model — it's a sign of covalence between the soft sulfide ion and d-metal cations.
Pyrite Structure
Compounds containing discrete S₂²⁻ ions adopt the pyrite or marcasite structure. The mineral iron pyrites ("fool's gold") has the formula FeS₂ and consists of Fe²⁺ and discrete S₂²⁻ anions in a rock-salt structure.
Group 12/16 Semiconductors
Sulfides, selenides, and tellurides of Group 12 elements form industrially important semiconductors (CdS, CdSe, CdTe, ZnSe) used in optoelectronic applications such as solar cells and light-emitting diodes.
16.12 Oxides of Chalcogens
(a) Sulfur Oxides and Oxohalides
SO₂ and SO₃ are both Lewis acids, with SO₃ being much stronger and harder. SO₂ is manufactured on a large scale by:
The main components of acid rain are nitric and sulfuric acids produced by interaction of the oxides with hydroxyl radicals:
Sulfuric and nitric acid particles are a major health threat, associated with increased mortality from pulmonary and heart disease.
(b) Oxides of Selenium and Tellurium
Selenium trioxide is thermodynamically less stable than the dioxide (unlike SO₃ or TeO₃):
| Compound | ΔfH° /kJ mol⁻¹ |
|---|---|
| SO₂ | −297 |
| SO₃ | −432 |
| SeO₂ | −230 |
| SeO₃ | −184 |
| TeO₂ | −325 |
| TeO₃ | −348 |
(c) Chalcogen Oxohalides
The most important oxohalides are thionyl dihalides (OSX₂) and sulfuryl dihalides (O₂SX₂).
16.13 Oxoacids of Sulfur
Sulfur forms many oxoacids existing in aqueous solution or as solid salts of oxoanions.
Sulfur Oxoanions
| Oxidation Number | Formula | Name | Remarks |
|---|---|---|---|
| +4 | SO₃²⁻ | Sulfite | Basic, reducing agent |
| +6 | SO₄²⁻ | Sulfate | Weakly basic |
| +2 | S₂O₃²⁻ | Thiosulfate | Moderately strong reducing agent |
| +3 | S₂O₄²⁻ | Dithionite | Strong reducing agent |
| +4 | S₂O₅²⁻ | Disulfite | — |
| +5 | S₂O₆²⁻ | Dithionate | Resists oxidation and reduction |
(b) Sulfuric Acid
H₂SO₄ is a dense viscous liquid that dissolves in water in a highly exothermic reaction:
It is a strong Brønsted acid (pKa1 = −2) and undergoes extensive autoprotolysis:
Contact Process
Concentrated sulfuric acid is manufactured by the contact process:
- Oxidation of sulfur to SO₂
- Oxidation of SO₂ to SO₃ over V₂O₅ catalyst
- Absorption of SO₃ in sulfuric acid
(d) Thiosulfuric Acid
The thiosulfate ion is a moderately strong reducing agent:
The reaction with iodine is the basis of iodometric titrations:
(e) Peroxodisulfuric Acids
Peroxodisulfate salts are strong oxidizing agents:
16.14 Polyanions of Sulfur, Selenium, and Tellurium
Many polysulfides of electropositive elements contain Sn²⁻ ions where n = 2−6. Typical examples are Na₂S₂, BaS₂, Na₄S₄, K₂S₄, and Cs₂S₆.
Structures
- Polysulfides (Sn²⁻) — chains up to n = 6
- Polyselenides — chains up to Se₉²⁻; larger ones form rings (e.g., Se₁₁²⁻)
- Polytellurides — chains and bicyclic structures (e.g., Te₇²⁻, Te₈²⁻)
The radical anion S₃⁻ occurs in the blue pigment ultramarine, embedded in pores of an aluminosilicate structure with Na⁺.
16.15 Polycations of Sulfur, Selenium, and Tellurium
Because these cations are oxidizing agents and Lewis acids, preparative conditions are quite different from those used to synthesize polyanions. For example:
Square-Planar E₄²⁺ Ions
Sulfur, Se, and Te each form ions of the type E₄²⁺ with square-planar (D4h) structures. In the MO model, six electrons fill the a2u and eg orbitals, leaving the antibonding b2u orbital vacant.
S₈²⁺ Ion
Oxidation of S₈ to S₈²⁺ brings about formation of an additional 2c,2e bond. X-ray structure shows transannular bonds are long (283 pm) compared with other bonds (204 pm).
16.16 Sulfur–Nitrogen Compounds
Tetrasulfurtetranitride (S₄N₄)
The oldest known and easiest to prepare sulfur-nitrogen compound. It is a pale yellow-orange solid made by:
S₄N₄ is endergonic (ΔfG° = +536 kJ mol⁻¹) and may decompose explosively. The molecule is an eight-membered ring with four N atoms in a plane and bridged by S atoms projecting above and below.
Disulfurdinitride (S₂N₂)
Formed when S₄N₄ vapour is passed over hot silver wool. It is even more sensitive than its precursor and explodes above room temperature.
Polythiazyl (SN)n
When S₂N₂ stands at 0°C for several days, it transforms into a bronze-coloured zigzag polymer which:
- Is much more stable than S₂N₂ (doesn't explode until 240°C)
- Exhibits metallic conductivity along the chain axis
- Becomes superconducting below 0.3 K
This was the first superconductor with no metal constituents. Halogenated derivatives have even higher conductivity.
Question: Predict whether S₂N₂ could be described as aromatic.
Answer: Each S atom has 6 valence electrons. With one lone pair and two bonds to N atoms, each S has 2 electrons for π bonding. Each N atom has 5 valence electrons; with one lone pair and two bonds to S, each N has 1 electron for π bonding. Total: 6 electrons for π bonding → S₂N₂ is aromatic (2n+2 rule with n=2).
Polysulfide polymers (SN)n have been found to be very effective for visualization of latent fingerprints. When S₂N₂ contacts a fingerprint, polymerization to blue-black (SN)n is induced and the print is exposed.
This provides an inexpensive, nondestructive, solvent-free imaging technique for forensic applications.
Exercises
Further Reading
- J.S. Thayer, "Relativistic effects and the chemistry of the heaviest main-group elements", J. Chem. Educ., 2005, 82, 1721.
- R.B. King, Inorganic chemistry of the main group elements. John Wiley & Sons (1994).
- D.M.P. Mingos, Essential trends in inorganic chemistry. Oxford University Press (1998).
- P. Ball, H₂O: a biography of water. Phoenix (2004).
- R. Steudel, Elemental sulfur and sulfur-rich compounds. Springer-Verlag (2003).
- N.N. Greenwood and A. Earnshaw, Chemistry of the elements. Butterworth-Heinemann (1997).
- P.R. Ogilvy, "Singlet oxygen: there is indeed something new under the sun", Chem. Soc. Rev., 2010, 39, 3181.