Lecture 5c. Structure of the elements, continued P, S and I.

Allotropes of phosphorus

Elemental phosphorus exists in a number of different allotropes.

allotropes of phosphorus

White phosphorus
The most important form of elemental phosphorus from the perspective of applications and the chemical literature is white phosphorus. It consists of tetrahedral P4 molecules, in which each atom is bound to the other three atoms by a single bond. This P4 tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of 800 °C when it starts decomposing to P2 molecules. Solid white phosphorus exists in two forms. At low-temperatures, the β form is stable. At high-temperatures the α form is predominant. These forms differ in terms of the relative orientations of the constituent P4 tetrahedra.

The history of the match is linked to the discovery of the allotropes of phosphorus.

Allotropes of Phosphorus



violet - Hittorf

Phosphorene is an allotrope of phosphorus normally used to designate a single layer of black phosphorus that may be somewhat flattened. Conceptually the structure is similar to the carbon-based graphene, hence the name phosphorene. However phosphorene is a semiconductor, unlike graphene which is a semimetal. Recently a sample that was about 20 layers thick was shown to demonstrate high-speed data communication on nanoscale optical circuits.

White phosphorus is the most reactive, the least stable, the most volatile, the least dense, and the most toxic of the allotropes. White phosphorus gradually changes to red phosphorus. This transformation is accelerated by light and heat, and samples of white phosphorus almost always contain some red phosphorus and accordingly appear yellow. For this reason, white phosphorus that is aged or otherwise impure is sometimes called yellow phosphorus. White phosphorus glows in the dark (when exposed to oxygen) with a very faint tinge of green and blue, is highly flammable and pyrophoric (self-igniting) upon contact with air and is toxic (causing severe liver damage on ingestion). Owing to its pyrophoricity, white phosphorus has been used as an additive in napalm. The odour of combustion of this form has a characteristic garlic smell, and samples are commonly coated with white "phosphorus pentoxide", which consists of P4O10 tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is insoluble in water but soluble in carbon disulfide.

Red phosphorus
In 1847 Anton von Schrotter found that sunlight changed white/yellow into red phosphorus, even when moisture and atmospheric oxygen were rigorously excluded. The red product was separated from the residual yellow phosphorus by treatment with carbon disulfide. Red phosphorus was also prepared from the yellow variety by heating it to about 250 °C. in an inert gas. Heating to higher temperatures reconverted the red modification to the yellow one.
Red phosphorus exists as an amorphous network and does not ignite in air at temperatures below 240 °C.

Violet phosphorus
In 1865, Johann Hittorf heated red phosphorus in a sealed tube at 530 °C. The upper part of the tube was kept at 444 °C. Brilliant opaque monoclinic, or rhombohedral, crystals sublimed.
This form is sometimes known as "Hittorf's phosphorus" (or violet or α-metallic phosphorus).

Black phosphorus
Black phosphorus is the thermodynamically stable form of phosphorus at room temperature and pressure. It is obtained by heating white phosphorus under high pressures (12,000 atmospheres). In appearance, properties and structure it is similar to graphite, being black and flaky, a conductor of electricity, and having puckered sheets of linked atoms.
Black phosphorus has an orthorhombic structure and is the least reactive allotrope: a result of its lattice of interlinked six-membered rings. Each atom is bonded to three other atoms.

Allotropes of sulfur

No other element forms more solid allotropes than sulfur. At present, about 30 well characterized sulfur allotropes are known of which the most common form found in nature is the greenish-yellow orthorhombic α-sulfur, containing puckered rings of S8.


When pure it has a greenish-yellow colour (traces of cyclo-S7 in commercially available samples make it appear yellower). It is practically insoluble in water and is a good electrical insulator with poor thermal conductivity. It is quite soluble in carbon disulfide: 35.5 g/100 g solvent at 25 °C. It has a rhombohedral crystal structure. This is the predominant form found in "flowers of sulfur", "roll sulfur" and "milk of sulfur". It contains S8 puckered rings, alternatively called a crown shape. The S-S bond lengths are all 206 pm and the S-S-S angles are 108° with a dihedral angle of 98°. At 95.3 °C, α-sulfur converts to β-sulfur.


This is a yellow solid with a monoclinic crystal form and is less dense than α-sulfur. Like the α- form it contains puckered S8 rings and only differs from it in the way the rings are packed in the crystal. It is unusual because it is only stable above 95.3 °C, below this it converts to α-sulfur. It can be prepared by crystallising at 100 °C and cooling rapidly to slow down formation of α-sulfur. It has a melting point of about 120 °C and decomposes at around this temperature.


This form, first prepared by F.W Muthmann in 1890, is sometimes called "nacreous sulfur" or "mother of pearl sulfur" because of its appearance. It crystallises in pale yellow monoclinic needles. It contains puckered S8 rings like α-sulfur and β-sulfur and only differs from them in the way that these rings are packed. It is the densest form of the three. It can be prepared by slowly cooling molten sulfur that has been heated above 150 °C or by chilling solutions of sulfur in carbon disulfide, ethyl alcohol or hydrocarbons. It is found in nature as the mineral rosickyite.

Some allotropes of Sulfur

S6 - cyclohexasulfur


S12 - cyclododecasulfur

S6 - cyclo-hexasulfur
This was first prepared by M.R. Engel in 1891 who reacted HCl with thiosulfate, HS2O3-. Cyclo-S6 is orange-red and forms rhombohedral crystals. It is called ρ-sulfur, ε-sulfur, Engel's sulfur and Aten's sulfur. Another method of preparation involves reacting a polysulfane with sulfur monochloride:
H2S4 + S2Cl2 → cyclo-S6 + 2 HCl (dilute solution in diethyl ether)
The sulfur ring in cyclo-S6 has a "chair" conformation, reminiscent of the chair form of cyclohexane. All of the sulfur atoms are equivalent.

Thermodynamically, S12 is the second most stable sulfur ring after S8. Therefore, S12 is formed in many chemical reactions in which elemental sulfur is a product. In addition, S12 is a component of liquid sulfur at all temperatures. The same holds for S18 and S20 which are often formed together with S12. Its structure can be visualised as having sulfur atoms in three parallel planes, 3 in the top, 6 in the middle and three in the bottom.

Liquid sulfur after equilibration contains sulfur homocycles of all sizes and some of these can be isolated by quenching, extraction, fractional precipitation and crystallization depending on their differing solubilities.

Cyclo-S12 can be prepared by heating elemental sulfur to about 200 °C for 5-10 min and then allowing the mixture to cool to 140-160 °C within about 15 min. Once the melt has become less viscous, it is poured in as thin a stream as possible into liquid nitrogen in order to quench the equilibrium. Recrystallization of the yellow powder from CS2 allows the isolation of an adduct which slowly loses the solvent to give the cyclo-dodecasulfur.

Note that both B and S form stable E12 species but the structures (and coordination numbers) are quite different.


Iodine was discovered by French chemist Bernard Courtois in 1811. His father was a manufacturer of saltpeter (a vital part of gunpowder) and at the time of the French Napoleonic Wars, saltpeter was in great demand. Saltpeter produced from French niter beds required sodium carbonate, which could be isolated from seaweed collected on the coasts of Normandy and Brittany. To isolate the sodium carbonate, the seaweed was burned and the ash washed with water. The remaining waste was destroyed by adding sulfuric acid. Courtois once added excessive sulfuric acid and a cloud of purple vapour rose. He noted that the vapour crystallized on cold surfaces, making dark crystals. Courtois suspected that this was a new element but lacked the funds to pursue it further.

Samples of the material reached Humphry Davy and Joseph Louis Gay-Lussac and in early December 1813 both claimed that they had identified a new element. Arguments erupted between them over who had identified iodine first, but both scientists acknowledged Courtois as the first to isolate the element.

Iodine is found on Earth mainly as the highly water-soluble iodide ion I-, concentrated in oceans and brine pools. Like the other halogens, free iodine occurs mainly as a diatomic molecule I2. In the universe and on Earth, iodine's high atomic number makes it a relatively rare element. However, its presence in ocean water has given it a role in biology. It is the heaviest essential element widely utilized by life in biological functions.

Under standard conditions, iodine is a bluish-black solid that sublimes to form a noxious violet-pink gas. It melts at 113.7 °C (386.85 K) and forms compounds with many elements but is less reactive than the other halogens, and has some metallic light reflectance.

Elemental iodine is slightly soluble in water, with one gram dissolving in 3450 ml at 20 °C and 1280 ml at 50 °C; potassium iodide may be added to increase solubility via formation of triiodide ions (I3-). Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility.

Iodine normally exists as a diatomic molecule with an I-I bond length of 270 pm, one of the longest single bonds known. The I2 molecules tend to interact via weak London dispersion forces, and this interaction is responsible for the higher melting point compared to more compact halogens, which are also diatomic. Since the atomic size of iodine is larger, its melting point is higher.

The I-I bond is relatively weak, with a bond dissociation energy of 151 kJmol-1, and most bonds to iodine are weaker than for the lighter halides. One consequence of this weak bonding is the relatively high tendency of I2 molecules to dissociate into atomic iodine.
orthorhombic structure of I2
a= 0.72701, b= 0.97934, c= 0.47900 nm

The halogens, Cl2, Br2, and I2 adopt similar orthorhombic structures in which diatomic molecules lie in layers:
Cl a= 0.624 b= 0.826 c= 0.448 nm
Br a= 0.667 b= 0.872 c= 0.448 nm
I a= 0.72701, b= 0.97934, c= 0.47900 nm

Return to the course outline or move on to Lecture 6: Acids, Bases and Solvent Systems.


Much of the information in these course notes has been sourced from Wikipedia under the Creative Commons License.
'Inorganic Chemistry' - C. Housecroft and A.G. Sharpe, Prentice Hall, 4th Ed., 2012, ISBN13: 978-0273742753, pps 24-27, 43-50, 172-176, 552-558, 299-301, 207-212
'Basic Inorganic Chemistry' - F.A. Cotton, G. Wilkinson and P.L. Gaus, John Wiley and Sons, Inc. 3rd Ed., 1994.
'Introduction to Modern Inorganic Chemistry' - K.M. Mackay, R.A. Mackay and W. Henderson, International Textbook Company, 5th Ed., 1996.

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