VIB GROUP (16) ELEMENTS
ns2np4: O, S, Se, Te, Po
OXYGEN
Oxygen compounds of all elements except He, Ne, and Ar known. Molecular O2 reacts, at room temperature or on heating, with all other elements except the halogens, noble metals, and noble gases.
Since O is in the 2nd short period, its chemistry is based on completing the octet. This can be achieved in more than one way:
1. e- gain
oxide O2-
2. Formation of 2 single covalent bonds in AB2E2systems (H2O) 3. Formation of a double bond in ABE2 systems (ketones, R2C=O)
4. Formation of a single bond + e- gain in ABE3 systems (OH-)
5. Formation of 3 covalent bonds in pyramidal AB3E systems (H3O+) 6. Rare: formation of 4 covalent bonds such as in basic Be acetate Be4O(CH3CO2)6 (Zn has an analogue compound):
Bond types in oxygen compound range from ionic to covalent, as do the physical properties of the oxides. Ionic oxides (alkaline- and earth-alkaline- metal oxides Completely covalent (molecular): CO2 Intermediate cases: oxides of B, Al, Si
IONIC OXIDES O2(g) O(g) O(g) + 2e- O2This 2-step process takes about 1000 kJ/mol. In addition energy must be provided for vaporization and ionization of the metal. All this endothermic energy is offset by the lattice energy provided for by a small and highly charged O2- ion. If the lattice energy (for various reasons) is not sufficient to make up all the endothermic energy needed, than the certain contribution of covalency is typically observed (BeO, SiO2, B2O3). COVALENT OXIDES Can use the ABxEy classification from valence theory. A=O, B is the other atom or molecular species, and E=lone e- pair p-p or to a lesser extent d-p double bonds are very important in many oxides' structures
ABE3 systems: O is sp3 hybridized, 1 bond, 3 lone pairs (OH-) AB2E2 systems: O is again sp3 hybridized, 2 bonds, 2 lone pairs; these systems are angular due to volume requirements of 2 lone pairs (H2O) AB3E systems: again sp3 hybridization, 3 bonds, 1 lone pair (H3O+); H3O+ is analogous to NH4+; however H4O2+ would not be possible (even with a lone pair on O) because of electrostatic repulsion of already positive ion towards H+. ABE2 systems: sp2 hybridization (ketones and aldehydes); O doubly bonded ACID-BASE PROPERTIES OF OXIDES oxides of metals are BASIC oxides of non-metals are ACIDIC AMPHOTERIC oxides
BASIC OXIDES
Oxide (O2-), peroxide (O22-) and superoxide (O2-) ions exist in solids, but not in solution due to the hydrolysis: O2- + H2O 2OHO22- + H2O HO2- + 2OHO2- + H2O O2 + HO2- + 2OH-
However, these ionic oxides are typically insoluble in water. They are usually soluble in acids. MgO(s) + 2H+(aq) Mg2+ + H2O
ACIDIC OXIDES
Frequently the term "acidic anhydride" is used for acidic oxides of covalent non-metal elements:
N2O5 + H2O 2H+ + 2NO3-
If acidic oxides are insoluble in water they will generally dissolve in bases: Sb2O5 + 2OH- + 5H2O 2Sb(OH)6-
AMPHOTERIC OXIDES
Amphoteric- acid behavior towards bases; basic behavior towards acids:
ZnO(s) + 2H+(aq) Zn2+ + H2O ZnO + 2OH- Zn(OH)42-
OTHER OXIDES
There are oxides that are relatively inert, and will not dissolve in neither acids nor bases (N2O, CO, PbO2, MnO2). When they do react, it is a redox reaction: MnO2(s) + 4HCl Mn2+ + 2Cl- + Cl2 + 2H2O
OXYGEN ALLOTROPES: O2 AND O3 (OZONE)
O2: paramagnetic, : - : : :
ITS GROUND STATE IS TRIPLET
Unpaired electrons in degenerate orbitals can have the same spin, so the total spin S of the molecule is 1. This is known as a triplet configuration because the spin has three possible alignments in an external magnetic field.
orbitals: Because of these unpaired electrons both gaseous and liquid O2 are pale blue.
SINGLET O2, molecular oxygen at higher energy level, in which all the electron spins are paired, In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight. It is also produced by the immune system as a source of active oxygen.
O3 OZONE
Ozone is a poisonous gas with a distinct, sharp odor. It is thermodynamically unstable compared to O2. It is formed continuously in the upper atmosphere of the Earth by short-wave ultraviolet (UV) radiation, and also functions as a shield against UV radiation reaching the ground.
It is powerful oxidizer (even more so than O2).
From Wikipedia:
A newly discovered allotrope of oxygen, tetraoxygen (O4), is a deep red solid that is created by pressurizing O2 to the order of 20 GPa. Its properties are being studied for use in rocket fuels and similar applications, as it is a much more powerful oxidizer than either O2 or O3.[6][7] When tetraoxygen is subjected to a pressure of 96 GPa, it becomes metallic, similarly to hydrogen, and becomes more similar to the heavier chalcogens, such as tellurium and polonium, both of which show significant metallic character.
COMPOUNDS OF OXYGEN
HYDROGEN PEROXIDE, H2O2 H O-O H Resembles water in many physical properties. It is even more highly associated by H-bonding (denser and higher BP). In dilute solutions more acidic than water. In spite of the high dielectric constant can not be used as solvent due to strong OXIDIZING properties. Actually it is a strong oxidizer in either basic or acidic solutions. It can behave as reducing agent toward very strong oxidizers, such as permanganate ion, MnO4-.
PEROXIDES AND SUPEROXIDES
derived formally from O22- and O2-
IONIC PEROXIDES: with alkali metals, Ca, Sr, Ba
In water or dilute acid solutions they produce H2O2. They are powerful oxidizing agents.
IONIC SUPEROXIDES: MO2, where M= K, Rb, Cs; yellow-orange crystalline solids. O2- has 1 unpaired e-, and superoxides are also powerful oxidizng agents.
Alkaline-earth, Zn, Cs superoxides occur only in small concentrations as solid solutions in peroxides.
DIOXYGEN AS A LIGAND
Oxidations: redox reactions, oxygen gets reduced, the other partner in the reaction gets oxidized Oxygenations: instead of acting as an oxidant, O2 acts as ligand, retains its identity; often reversible reaction REVERSIBLE OXYGENATION ESSENTIAL FOR LIFE: hemoglobin and mioglobin contain 1:1 Fe-O2 complex that carries the intact O2 through the tissues. Oxygenations can be END-ON and SIDE-ON:
AQUA LIGANDS H2O molecules can be weakly bound (and very rapidly substituted (e.g. with alkali-metal cations)
in other cases, such as [Cr(H2O)6]3+ H2O ligands can be very firmly bound we already mentioned cases where these complexes are weakly acidic (group IIIB metals) HYDROXO LIGANDS (OH)
Many complexes are known. 2 possibilities: terminal and bridging ligands (look up example in Fig. 18-3 on p. 446 in Cotton)
some metal ions, such as Co(III) and Al(III) will have structurally different aqua (octahedral) and hydroxo (tetrahedral) complexes
OXO LIGANDS A double O=M bond of the type OpMd exists in number of molecules or molecular fragments: vanadyl (V=O), uranyl (U=O), permanganate MnO4-, and OsO4. Metal oxo complexes are most stable when metal is in high oxidation state by contrast to non-metals, where low oxidations states are preferred (CO2 and SO2).
OXYGEN FLUORIDES called fluorides, since F is more electronegative then O; possible application as rocket fuel oxidizers: OF2 and O2F2; the latter is stable only up to -50 C.
SULFUR, SELENIUM, TELLURIUM, AND POLONIUM
S occurs widely in nature as an element, H2S, SO2, in the form of metal sulfide ores or sulfates. Se and Te are less abundant, but commonly found as selenides and telurides alongside sulfide ores.
Po occurs in U and Th minerals as a product of radioactive decay.
Elements S through Po have fairly low MP's ( 200400 C), coming in various forms.
S has a profusion of allotropes, that can contain multiple atom rings (up to 20 atoms) or in chains (plastic or catena sulfur). The most stable form is orthorhombic S8.
The most stable form of Se is the gray, metal-like spiral chains, with weak interaction of the metallic nature between different atoms of neighboring chains. These chains are photoconductive, which is very useful property for photoelectric devices.
Te is isomorphous with grey Se, but its color is silvery white. Based on the resistivity it is classified as semimetalic.
The resistivities of S, Se, and Te all have negative temperature coefficients, typical for non-metals.
CHEMISTRY OF S, Se, Te, and Po
1. The electronegativities fall in the VIB group from O through Po, so does the ionic character of their compounds. Also the H-bonding becomes increasingly less important. H2S is very different from H2O.
2. d-p bonding is a lot more important for S than p-p bonding. 3. The valence of S, Se, Te, and Po are not confined to 2, utilizing d orbitals to form more than 4 bonds to other elements (SF6 and Te(OH)6). 4. S has a strong tendency for catenation, exceeded only by C. There are no known analogues in O, Se or Te chemistry. Examples are polysufide ions, Sn2-, polythionate ions, [O3S-Sn-SO3]2-, and compounds of type XSnX (X=H, Cl, CN, or NR2). S-S bonds occur in varieties of compounds and S-S- bridges are especially important in certain enzymes and other proteins. The changes in properties of compounds going from S to Po are a function of increasing sizes of atoms and decreasing electronegativity. Some typical trends are: 1. Decreasing stability of hydrides H2E. 2. Increasing tendency to form complex ions such as SeBr62-.
3. Appearance of metallic properties for Te and Po atoms. Oxides MO2 are becoming increasingly ionic and basic, reacting with HCl to give chlorides. REACTIONS AND COMPOUNDS Salts of cations M42+, M82+, and M162+ can be synthesized by selective oxidation of the elements with SbF5 or AsF5 in liquid HF or by reactions in molten AlCl3: S8 + 3 SbF5 S82+ + 2 SbF6- + SbF3 7Te + TeCl4 + 4 AlCl3 2Te42+ + 4 AlCl4The M42+ ions are square, likely most featuring a 6 -electron quazi-aromatic system. The M82+, and M162+ have a ring and fused-ring structures, respectively:
These cations are highly colored and atoms have fractional oxidation state.
VULCANIZATION an extremely important industrial process addition of S to double bonds in natural and synthetic rubbers, creates bridges between C chains and strengthens the rubber.
HYDRIDES (EH2)
Extremely poisonous and malodorous. Their thermal stabilities and bond strengths fall down the group, and the acidities increase.
H2S dissolves in water to give mostly protons and HSions, since the second dissociation constant is much smaller than the first. S2- ions occur only in very alkaline solutions. SULFANES SSSSSS- chains. Anions S42- and S52- are very good chelating ligands: H2S2 to H2S6 contain SS- up to
HALIDES AND OXOHALIDES
Direct fluorination of S8 yields SF6 with traces of S2F10 and SF4. SF6 is very resistant to chemical attack (due to high SF bond strength, coordinative saturation and steric hindrance at S atom). Its inertness, molecular weight and high dielectric strength allow it to be used in high voltage generators as an insulator. This is an example where a compound is inert due to the kinetic, rather than thermodynamic factors.
Thionyl chloride, SOCl2 is made by: SO2 + PCl5 SOCl2 + 2HCl SOCl2 is a weak Lewis base and mild chlorinating agent.
OXIDES AND OXOACIDS SO2 and H2SO3 : O : : O: : :S
+
-
S actually has an extended octet (d orbitals available), considerable p-d bonding, planar structure
SO2 can act as a Lewis base AND Lewis acid. Can bind in complexes through S or through O.
SO2 is also a reducing agent, since S is not in its highest oxidation state. SO2 is gas soluble in water to form solutions which were in the past referred to as SULUROUS ACID. The latter actually does not exist in solutions, but it has a form of hydrate, SO2x~7H2O, which still dissociates to form HSO3- and H3O+. However, the alkali-metal salts containing HSO3- and SO32- ions do exist, and they are called bisufites and sulfites respectively. On heating 2 moles of bisulfites give up 1 mole of H2O to yield PYROSULFITES M2S2O5.
SeO2 is a solid which has a chain structure, it is very volatile and in gas phase it has similar structure to SO2. TeO2 is also a solid and it crystallizes in rutile crystal lattice, with very strong bonds and it is not volatile at all.
SO3 and H2SO4 Catalyzed reaction (thermodynamically favorable, but very slow: 2SO2 + O2 2SO3 planar, triangular structure, involving both p-d and p-p S-O bonding; forms polymers in the solid :O= S
:
++
:O: :O:
:
Anhydrid of H2SO4. H2SO4 has tetrahedral structure, its salts are bisulfates (HSO4-) and sulfates (SO42-). Selenic acid H2SeO4 is similar to sulfuric acid, but less stable. Telluric acid Te(OH)6 is very different. As we discussed, it is a very weak base. Most tellurates contain TeO6 octahedra.
There is a number of other S acids containing more than 1 S atom, obtained either by redox reactions or dehydratation.
:
-
GROUP VIIB (17) ELEMENTS (HALOGENS)
ns2np5: F, Cl, Br, I, At
With the exception of He, Ne, and Ar, all of the elements in the periodic table form ionic or covalent halides. They are among the most important and most common compounds. There is also an extensive organic-halogen chemistry.
Fluorine occurs in several different minerals, and it is more abundant than Cl. F2 is most chemically reactive of all the elements and combines directly at ordinary or elevated temperatures with all elements except O2, He, Ne, and Kr. It attacks most of the compounds, inorganic or organic, often inflaming them. This great reactivity of F2 is partly due to its very low dissociation energy, which again is caused by repulsion between the non-bonding electrons. The other factor that influences its reactivity is the exothermicity of its reactions.
Chlorine occurs in the rock salt and many other minerals, in the seawater and many prehistoric deposits formed by drying out of huge bodies of water. It is a green-yellow gas, somewhat less reactive than F2, it is moderately soluble in water by the reaction:
Cl2 + H2O HCl + HOCl (hypochloric acid)
Bromine is a dark red liquid, moderately soluble in water and miscible with polar solvents. Br is found in nature in the form of bromides alongside chlorides.
Iodine occurs as iodate in Chile saltpeter (deposits naturally made out of bird droppings, containing mostly nitrites), and as iodine in sea water. I2 is a black solid with slight metallic luster, which sublimes without melting.
With non-polar solvents it forms purple solutions, while solutions in polar solvents, liquid SO2 and unsaturated hydrocarbons are pinkish-brown (due to formation of weak complexes of I2, so called "charge transfer complexes". The bonding energy results from a partial transfer of charge from solvent to I2 (I2-S+). I2 forms a blue complex with starch, forming linear I5ions in channels of polysaccharide amylose.
Astatine has been identified as a short-lived product of radioactive decay of U and Th. Based on limited experimental evidence it is believed its chemical behavior is in line with halogen behavior in general considering its place at the bottom of the group.
CHEMICAL BEHAVIOR
The halogen atoms are only 1 e- short of the noble gas configuration, so the most dominant form of chemical behavior is a formation of X- or a single covalent bond.
Halogen chemistry is completely non-metallic. The changes in behavior with increasing size are progressive, however this is the most cohesive group in the periodic table.
Apart from monovalent state, all halogens with exception of F, form compounds in which they maintain oxidation states of +1, +3, +4, +5, and +7, by loosing the appropriate number of e-. These compounds are either interhalogens or oxo compounds.
No evidence exist of cationic behavior of the type X+, however, Br2+, I2+, Cl3+, and Br3+, in addition to several larger I cations, are known.
When halogen forms a bond to another halogen, the bond will be polar with a positive charge on heavier halogen (electronegativities drop down the group!).
COMPOUNDS IONIC HALIDES
MOLECULAR (COVALENT) HALIDES
formed by most electronegative elements and metals in high oxidation states; they are gases, liquids, or volatile solids held together by van der Waals forces
important structural feature: formation of halide bridges, typically 2 bridges, but compounds with 1 or 3 bridges are known: X M M X molecular halides EXn are very easily hydrolyzed:
the occurrence of F as a substituent in many compounds can dramatically emphasize or change the chemical behavior of that molecular species, e.g.: 1. F3CCOOH is extremely strong acid, due to cumulative electronegativity of 3 F atoms
2. (CF3)3N and F3N have literally no basicity OXIDES AND OXYACIDS
OF2 and O2F2; extremely potent oxidizing and fluorinating agents. Studied as potential rocket fuels.
Oxidation state +1 Cl2O (powerful oxidant, used for bleaching), HClO (hypochloric acid), its salts hypochlorites (ClO-) are stable in alkaline solutions). Formed by introduction of Cl2 into water: Cl2 + H2O H+ + Cl- + HClO
Br and I analogues are progressively less stable.
Oxidation states +4 and +3 very explosive ClO2 (+4 oxidation state), also used for bleaching, it is related to: HClO2 (chlorous acid) , weak acid and not very stable, stable alkali salts-chlorites (MClO2) If we introduce ClO2 into alkali solution it disproportionates to hypochlorite and chlorate (oxidation state +5): 2ClO2 + 2OH- ClO- + ClO3- + H2O
No Br or I analogues are known.
Oxidation state +5
Halic acids (HXO3): chloric, bromic, iodic, and their salts chlorates, bromates, and iodates All very strong acids and powerful oxidizing agents.
Ions XO3- all have pyramidal structure:
++
-
- O
Cl O - O
Note: formal charge is +2, and oxidation state is +5, total ion charge -1; Cl lost 2 e- and the other 3 are engaged in Cl-O bonding.
:
Oxidation state +7 Perchloric acid HClO4 is the strongest inorganic acid. Its salts are perchlorates. HBrO4 is very unstable, but its salts, perbromates are known, they are extremely powerful oxidants.
Per-iodic acid is known in the forms meta per-iodic HIO4 or ortho per-iodic H5IO6 (it has 2 extra water molecules). Periodates are known and resemble tellurates in their stoichiometries.
The structure of ClO4- is again tetrahedral: O+++
-O
Cl
OO-
Perchlorate ion is extremely stable, it is extremely weak base, it does NOT have any oxidative properties and it does not form complexes.
ORGANOFLUORINE COMPOUNDS
High C-F bond energy (C-F bond 486 kJ/mol, C-H bond 415 kJ/mol). They are relatively inert, due to inability of F atom to expand an octet which hinders formation of transition states by coordinating other molecular species. F easily replaces H atoms since it is relatively small in size and then shields C atom. C in C-F is effectively oxidized (C in C-H is reduced), so fluorocarbons are inert to oxidation. They also have low boiling points and low friction.
Chloroflurocarbons have been used as nontoxic, inert refrigerants, but they are being phased out since they are prone to releasing Cl radicals in the upper atmosphere, which were proven to contribute a great deal to the destruction of the ozone layer. At one point, completely fluorinated hydrocarbons were considered as a replacement.
Another important use: fire supressants
Haloalkane, or halogenoalkane, a group of chemical compounds consisting of alkanes with linked halogens Various gaseous fire suppression agents: Halon 1211 (bromochlorodifluoromethane, CF2ClBr) Halon 1301 (bromotrifluoromethane, CBrF3)
(Wikipedia)
Great fire suppressants, but again issue with ozone layer destruction!!!!