Croup 17 (VIIA) - The group 17 of the modern periodic table consists of Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and Astatine (At).

- These elements are known as halogens and have nonmetallic character.

- All Group 17 elements have the electron configuration ns2np5 in their outer shell.

- Because the halogen elements have seven valence electrons, they only require one additional electron to form a full octet. This characteristic makes them: + More reactive than other non-metal groups. + Excellent oxidizing agents, with F2 being the strongest oxidizing agent of all the elements.

Occurrence:

- The halogens exist, at room temperature, in all three states of matter: Solid: Iodine, Astatine Liquid: Bromine Gas: Fluorine, Chlorine

- Fluorine most commonly occurs in the minerals cryolite, fluorspar, and fluorapatite. 0.06% of the Earth’s crust is fluorine, which makes fluorine the 13th most common element in the crust.

The halogens never occurs as a free element in nature:

- Chlorine is combined chiefly with sodium as sodium chloride in common salt (NaCl) and the minerals carnallite and sylvite.

- The most electronegative group, electronegativity decreases down the group.

- Bromine can be found combined with other elements and in salt water and springs.

- Salt mines contain some Iodine along with sea water but the main source is sea kelp.

- Astatine is radioactive element occurs naturally from uranium and thorium decay. But, it occurs chiefly from man-made.

States:Oxidation

- Oxidation numbers are typically are typically -1.

- Fluorine exhibits only -1 oxidation state in its

compounds since it is the most electronegative atom and the absence of vacant d-orbitals in valence shell.

- However elements such as chlorine, bromine and iodine

also show +1, +3, +5 and +7 state. This higher oxidation state of chlorine, bromine and iodine is realized when these halogens are in combination with small and highly electronegative atoms of fluorine and oxygen.

- Also, these elements contain vacant d–orbitals in their valence shells.

Physical properties are summarized in the following table:

Halogen F Cl Br I At

Electron Affinity (KJ/mol) -328 -349 -324.6 -295 -270

Electronegativity 4 3 2.8 2.5 2.2

1st Ionization Energy (kJ/mol) 1681 1251 1140 1140 890

Melting Point (˚C) -220 -101 -7.2 114 302

Boiling Point (˚C) -188 -35 58.8 184 337

- Fluorine (smallest size): is the most electronegative

highest ionization energy. And

- Astatine (largest size): is lowest.

- The London interactions between the molecules increase down the group because of increasing size so:

F2 and Cl2 are gases, Br2 is a liquid, and I2 is a solid.

- The high reactivity of F2 is because: of the weakness of the F-F bond;

+ The small size of the fluorine atom brings electron pairs closer when F-F bonds are formed. Repulsions between these pairs result in weaker bonding.

- These are compounds containing two or more different halogens. May be Diatomic: ClF or

Polyatornic: ClF3, BrF5, or IF7

- Selected Neutral Compounds:

Interhalogens

- The effect of the size of the central atom can be seen, with iodine the only element able to have up to seven fluorine atoms in a neutral molecule, whereas chlorine and bromine have a maximum of five fluorines.

Compounds Oxidation state of the central atom

ClF, BrF, IF, BrCl, ICl, IBr +1

ClF3, BrF3, IF3, I2Cl6 +3

ClF5, BrF5, IF5 +5

IF7 +7

1/ Direct reaction of the elements (the required ratio of halogens used):

C12 + F2 → 2 ClF

I2 + 5 F2 → 2IF5

2/ Reaction of halogens with metal halides or other halogenating agents:

I2 + 3XeF2 → 2IF3 +3Xe

I2 + AgF → IF + AgI

3/ Using Interhalogens as intermediates in the synthesis of other interhalogens:

ClF + F2 → ClF3 T = 200°C to 300°C

ClF3 + F2 → ClF5 0 oC

Preperation of Neutral interhalogens

Croup 18 (VIIIA)

The elements in Group 18 (VIIIA), long named the "inert or "rare" gases. They are now known to have an interesting, although limited, chemistry, and they are rather abundant. Helium, for example, is the second most abundant element in the universe, and argon is the third most abundant component of dry air.

- The Group 18 consists of elements Helium (He), Neon (Ne) Argon (Ar), Krypton (Kr), Xenon (Xe) and Radon (Rn).

- These elements have a particular name Noble gases. - Except helium, electron configurations: ns2p6 - Their closed shell electron configuration makes them

have a very low reactivity.

THE ELEMENTS

- The first experimental evidence for the noble gases was obtained by Henry Cavendish in 1766. In a series of experiments on air, he was able to remove nitrogen, oxygen and carbon dioxide from air by chemical means.

- Radon, was isolated as a nuclear decay product in 1902.

- In 1868, J. N. Locklear and E. Frankland proposed the existence of a new element named, appropriately, helium. - Lord Rayleigh and William Ramsay with Ramsay discovered a gas that might fit into the periodic table after the element chlorine. In 1895, they reported the details of their experiments and evidence for the element they had isolated, argon.

- Within 3 years, Ramay and M. W. Travers had isolated three additional elements by low-temperature distillation of liquid air, neon, krypton, and xenon. The last of the noble gases.

- Helium is rare on Earth, but it is the second most abundant element in the universe and is a major component of stars. Commercially, helium is obtained from natural gas. - The other noble gases, with the exception of radon, are present in small amounts in air and are commonly obtained by fractional distillation of liquid air. . - Argon, the least expensive noble gas, used for filling incandescent bulbs. - All isotopes of radon are radioactive; A potential cause of lung cancer.

Rn Xe Kr Ar Ne He Property

86 54 36 18 10 2 Atomic number, Z

202 161 116 84 24.5 ** Melting point, mp/K

211 165 120 87 27 4.2 Boiling point, bp /K

1037 1170 1351 1521 2081 2372 First ionization energy, kJ/mol

** Helium cannot be solidified under any conditions of temperature and pressure.

Some properties of Noble Gases

- For many years, these elements were known as the "inert" gases because they were believed to be totally unreactive because of the very stable "octet" valence electron configurations of their atoms.

- Attempts had been made to react xenon with elemental fluorine (the most reactive of the elements), but without success. However, in 1962, Neil Bartlett had observed that the compound PtF6 changed color on exposure to air.

- With D. H. Lohmann, he demonstrated that PtF6 was serving as a very strong oxidizing agent in this reaction and that the color change was due to the formation of a complex, noted the similarity of the ionization energies of xenon (1 169 kJ/mol) and O2 (1 175 kJ/mol) and repeated the experiment, reacting Xe with PtF6.

Chemistry

He observed a color change from the deep red of PtF6 to orange-yellow and reported the product as Xe+{PtF6}- the product of this reaction later proved to be a complex mixture of several xenon compounds.

- Bartlett's original reaction of xenon with PtF6 is now believed to proceed as follows:

Xe + 2PtF6 → XeF+(PtF6)- + PtF5 → XeF+(Pt2F11)-

- These were the first covalently bonded noble gas compounds to be synthesized, and their discovery began the study of the chemistry of the noble gases . In a matter of months, the compounds XeF, and XeF4 had been characterized.

- Other noble gas compounds soon followed 60 types of compounds of noble gas elements are now known, although the number is few in comparison with the other groups. The most known noble gas compounds of xenon and krypton.

- There is evidence for the formation of such radon compounds as RnF2, but the study of radon chemistry is difficult because of the high radioactivity.

- Recently, the first "stable" compound of argon, HArF, has been reported. Transient species containing helium and neon have been observed using mass spectrometry. However, most of the stable noble gas compounds are those of xenon with the highly electronegative elements F, O, and Cl; a few compounds have also been reported with Xe - N, Xe - C, and even Xe - transition metal bonds.

- Some of the compounds and ions of the noble gases are shown in Table:

- Several of these compounds and ions have interesting structures . For example, structures of the xenon fluorides have been interpreted on the basis of the VSEPR model (Figure).

XeF2 and XeF4 have structures entirely in accord with their VSEPR descriptions: XeF2 is linear (three lone pairs on Xe) and XeF4 is planar (two lone pairs). (see above figures)

- Krypton forms several species with fluorine, like the neutral KrF2. KrF2 exists in two forms in the solid. In the alpha form, shown in Figure. all molecules are parallel to each other, with eight molecules centered at the corners of the unit cell and a ninth centered in the cell.

Structure of XeF2

(Xe =yellow. F= green)

KrF2 (Kr=red, F= green)

(Xe2F11)2(NiF6)

(Xe=yellow, F=green Ni=blue

- Interest in using noble gas compounds as reagents in organic and inorganic synthesis is important because the byproduct of such reactions is often the noble gas itself. The xenon fluorides XeF2, XeF4, and XeF6 have been used as fluorinating agents for both organic and inorganic compounds. For example:

2SF4 + XeF4 → 2SF6 + Xe

- The oxides XeO3 and XeO4 are extremely explosive and must be handled under special care. XeO3 is a powerful oxidizing agent in aqueous solution. - The chemistry of krypton is much more limited, with fewer than a dozen compounds reported to date. The only neutral halide is KrF2. - The radioactivity of radon has made the study of its chemistry difficult; RnF2 and a few other compounds have been observed in few studies.