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Extract from chalcogen article 25/8/14

Atomic and physical

Chalcogens show similar patterns in electron configuration, especially in the outermost shells, where they all have the same number of valence electrons, resulting in similar trends in chemical behavior:

More information Z, Element ...

Z	Element	Electronic configuration	Electron Affinity (eV)^[1]	Ionisation potential^[2]	Electronegativity^[2]	Common oxidation states^[3]
8	oxygen	[He]2s² 2p⁴	1.461	13.61805	3.44	–2
16	sulfur	[Ne] 3s² 3p⁴	2.077	10.36001	2.58	-2, 2,4,6
34	selenium	[Ar] 3d¹⁰ 4s² 4p⁴	2.021	9.75239	2.55	-2, 2, 4, 6
52	tellurium	[Kr] 4d¹⁰ 5s² 5p⁴	1.971	9.0096	2.1	-2, 2, 4, 6
84	polonium	[Xe] 4f¹⁴ 5d¹⁰ 6s² 6p⁴	1.9	8.414	2.0	-2, 2, 4
116	livermorium	[Rn] 5f¹⁴ 6d¹⁰ 7s² 7p⁴ (predicted)^[4]		7.5 (predicted)^[5]		2, 4 (predicted)^[5]

More information Element, Melting point °C ...

Element	Melting point °C	Boiling point °C ^[6]	Density at STP (g/cm³)	Reference	State & color (STP) stable allotrope	structure stable allotrope	Band gap ^[7]	Other allotropes
Oxygen	−219	−183			gas , colorless	O₂	insulator	O₃
Sulfur	120	445	2.07	^[6]	solid, yellow,	S₈, puckered rings	insulator	many , ring structures and polymeric
Selenium	221	685	4.18	^[8]	solid grey metallic	Se_∝ helical chains	178	ring structures
Tellurium	450	988	6.25	^[9]	solid, grey metallic	Te_∝ helical chains	32.2
Polonium	254	962	9.2	^[6]	solid, metal	simple cubic	0
Livermorium	(435)^[10]	(812)^[10]	(12.9)^[5]

More information Element ...

Element
Oxygen
Sulfur
Selenium
Tellurium
Polonium

All chalcogens have six valence electrons. All of the solid, stable chalcogens are soft^[11] and do not conduct heat well.^[6] Electronegativity decreases towards the chalcogens with higher atomic numbers. Density, melting and boiling points, and atomic and ionic radii^[12] tend to increase towards the chalcogens with higher atomic numbers.^[6]

Trends in chalcogen chemistry

Greenwood _ Share tendency with group1 4, 15 elemnst to increasing metallic character with increasing atomic weight. O and s are insulators, Se and Te semiconductors and Po is a metal.

Te has some cationionic (basic ) caharcter and this is more pronounced in Po. Se is not apprciably attacked by dil. HCl Te dissolves to some extent in the presence of air Po dissolves readily . The structure and bonding of halides depends on both the electronegativity of the halogen and the oxidation state of the chalcogen, a ionic-covalent transition similar to that seen for group 15 . All fo the chalcogens combine directly with most elemnts Se, Te and Po combine less reqdily than O and S. The chalcogenides Se Te and Po form stable chalcogenides (M2-) with groups 1,2 and the lanthanides. compounds of Se, Te and Po with the more electronegative elemnts O,F,Cl have O states of 2,4 and 6. these compounds are less stable than corresponding compounds of S. Few analogues of the extensive range of Sulfur-nitrogen compounds. Decreasing thermal stability of hydrides with increasing atomic number. Se and Te share with s the propensity to catenate althogh to alesser extent. Diminution of stabilty of double bonds (e.g. to C, N and O) with increasing atomic number. CO2 iand to alesser extent CS2 are stable but CSe2 polymerises, SeCTE is unstable and CTe2 is unknown. SO2 is a non-linear molecule wheras SeO2 is a chain polymer, TeO2 has a layer or 3D structure. PoO2 adopts a typical ionic fluorite structure with 8 coordinate Po atoms. Double bonds are less readily formed the greater greater the electronegativity difference between them and the smaller the sum of their electronegativities. Se like other elemnts following the first transtion (3d series) shows a resistance to oxidation to its oxidation state VI. HNO3 readily oxidies S to H"SO4 wheras selenium only gibves H2SeO3. Dehydration of H2SO4 with P2O5 gives SO3 whears H2SeO4 produces SeO2 and O2. S forms many sulfones R2SO2 but relativley few selenoes are known.

SeO3 is unstable with respect to the dioxide ; wheras SO3 and TeO3 are not. SO3 in the solid is trimeric, SeO3 is tetrameric, wheras TeO3 has a structure similar to FeF3 with octahedral coorinated Te atoms sharing vertices to give a 3 dimensional structure. TeO3 is unaffected by water but is a strong ox agent when heated with metals. TMe6 is much more stable than TeMe4. HOWI (English)

Bonding Selenium has a lower tendency than S to form double bonds, p-p pi and p d pi bonds. Se is less likely to form elemnt elemnt bonds; H2Se2 is unstable unlike H2S2. Se2O52- has the structure O2Se-O-SeO2 wheras S2O52- is O2S-SO3. Tellurium forms double bonds less readily than selenium. The ability to form clusters where Te ah a coordination number >2 is greater than that of Selenium. H2Te is highly endothermic+99.6 kJ/mol. Te2I is an intercalation compound with I2 between layers of Te. TeO2 has highre coordination number of 4 than Se in SeO2 but lower than that of Po in PoO2 (8). H2TeO6 is actually Te(OH)6 unlike H2SO6 which is a hydrate H2SO4.H2O. Te forms t eTeO66- ion which is octahedral. (The first SeO6 compound Na12(SeO6)(SeO4)3 (in deVillanova 2007 along with SeO54-). Te forms solid TeO3. TMe6 is known

Hosecroft HSAB theory trends stability of complexes with chalcogen ions or ligands containing these donor atoms with hard ions shows the sequence O>> S>Se>Te wheras with soft ions O<<S >Se ~ Te.

Typical hard donors (O donor) include ROH, R2O, RO- Typical S donor ligands incudes RSH, R2S, RS-.

Carbonate

Decomposition stabilities

Basic carbonates

Basic carbonates are compounds where hydroxide or oxide ions are present along with carbonate ions. Sometimes the word basic is put before the metal sometimes after, for example basic copper carbonate or copper basic carbonate. Some of the more important ones industrially are bismuth basic carbonate, (BiO)₂CO₃·1.5H₂O; basic cobalt carbonate Co₅ (CO₃)₂ (OH)₆·H₂O; basic copper carbonate ??; basic lead carbonate (hydrocerussite) Pb₃ (CO₃)₂(OH)₂; basic magnesium carbonate Mg₅(CO₃)₄(OH)·5H₂O ; basic mercury carbonate Hg₄(CO₃) O₃ ; basic nickel carbonate Ni₃ CO₃ (OH)₄·4H₂O (4.5 H2O according to galwey and brown thermal decomposition of ionic solids) other formulae given elsewhere.

Coordination number

Crystalline substances

Determining the coordination number of an atom in a crystalline solid is sometimes difficult. Whereas in a molecule or polyatomic ion the bonded atoms can be counted, in the solid state it is not always clear which of the neighboring atoms should be included. The atoms that are included in the calculation of the coordination number define the vertices of what is called the coordination polyhedron. For example coordination numbers in the solid state of 4 and 6 are often associated with tetrahedral and the octahedral coordination polyhedra.^[13] Sometimes these coordination polyhedra are irregular, and it can bedifficult to determine the description to be given as polyhedra such as trigoanl bipyramd and the tetragonal pyramid can be interconverted by the displacelmnt of 4 vertices.^[14]

Metals

In metals where the structure is hexagonal close packed each atom is usually said to have 12 nearest neighbours and counting these gives a coordination number of 12. The ideal axial ratio of the unit cell in hexagonal close packing is 1.633 but this is only approximated to by real metals. Zinc metal has a distorted hexagonal close packed structure where the axial ratio of the unit cell is 1.856. In zinc there are there are only 6 nearest neighbours in the same close packed plane with six other, next nearest neighbours, equidistant, three in each of the close packed planes above and below. The 6 next nearest neighbours are generally only slightly more distant then the 6 nearest neighbours (291 pm compared to 266 pm^[15]) It is believed that these next nearest neighbours are significant and that the coordination number is therefore better described as 12 rather than 6.^[16] In the case of metals with the bcc structure there are 8 nearest neighbors with 6 more, approximately 15% more distant,^[17] and in this case the coordination number is often expressed as 14, or (8,6). In some complex cases such as white tin (β - tin), tin atoms have as near neighbours, 4 at 302 pm, 2 at 318 pm, and 4 at 377 pm. This sort of specification f coordination is often seen in inorganic chemistry^[18]^[19]

A number of different mathematical methods of unambiguously defining the coordination number have been developed which use the concepts of Voronoi polyhedra (sometimes called the domain of influence or Dirichlet domains) which restricting the atoms considered to be relevant to coordination by their distance from the central atom either by and in some methods these contributions are weighted giving rise to a fractional coordination number, the effective coordination number.

As a cautionary note has been suggested that the concept of coordination number implicitly assumes a central bonding force model and this assumption may not be valid in some crystalline substances such as metals.^[20]

Chemical compounds

For chemical compounds where there are always at least two different elements present the definition of a coordination number is further complicated. Take for example lithium iodide, LiI, which has the cubic rock salt (NaCl) structure with a close packed array of iodine atoms. Lithium iodide is essentially an ionic compound and the coordination numbers for both lithium and iodide ions are normally stated as 6. However the coordination number of an iodine atom could be considered to be 18 if, as well as the 6 lithium atoms the 12 iodine atoms with which it was also in contact were included.^[21]

Considerations like this have lead to the very broad definition adopted by the International Union of Crystallography, IUCR, which states that coordination number of an atom in a crystalline solid depends on the chemical bonding model and the way in which the coordination number is calculated.^[22]

Various mathematically derived methods have been suggested to determine an unambiguous coordination number in such cases.

References

[1]
Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, Florida: CRC Press. ISBN 0-8493-0487-3.
[2]
[3]
greenwood
[4]
Morss, Lester R; Edelstein, Norman M; Fuger, Jean, eds. (2006). The Chemistry of the Actinide and Transactinide Elements. Dordrecht, The Netherlands: Springer Science+Business Media. Bibcode:2011tcot.book.....M. doi:10.1007/978-94-007-0211-0. ISBN 978-94-007-0210-3. {{cite book}}: |journal= ignored (help)
[5]
Fricke, Burkhard (1975). "Superheavy elements: a prediction of their chemical and physical properties". Recent Impact of Physics on Inorganic Chemistry. Structure and Bonding. 21: 89–144. doi:10.1007/BFb0116498. ISBN 978-3-540-07109-9. Retrieved 4 October 2013.
[6]
Jackson, Mark (2002). Periodic Table Advanced. Bar Charts Inc. ISBN 978-1-57222-542-8.
[7]
greenwood
[8]
"Greenwood"
[9]
"Greenwood"
[10]
Bonchev, Danail; Kamenska, Verginia (1981). "Predicting the Properties of the 113–120 Transactinide Elements". J. Phys. Chem. 85 (9): 1177–1186. doi:10.1021/j150609a021.
[11]
Sampsonov, G.V. (1968). "Mohs Hardnesses of the Elements". IFI-Plenum. Retrieved November 25, 2013.{{cite web}}: CS1 maint: url-status (link)
[12]
"Visual Elements: Group 16". Rsc.org. Retrieved November 25, 2013.{{cite web}}: CS1 maint: url-status (link)
[13]
Muller
[14]
Muller
[15]
Wells
[16]
Mittemeijer
[17]
?Wells
[18]
Wells
[19]
Greenwood
[20]
Mittemeijer
[21]
Muller
[22]
IUCR