22 -- Part 2: Chemistry of the Main Group Elements II:
The shapes of the XeF2 and XeF4 molecule are easy to understand, but the shape of the XeF6 molecule is difficult to interpret.
These are capped trigonal and capped bipyramid.
The preferred arrangement depends on the exact conditions and the structures are expected to have nearly the same energy.
The capped octahedral structure of the XeF6 molecule is found in the gas phase.
The six fluorine atoms form a distorted octahedron, and the lone pair on xenon is directed toward the center of one of the triangular faces.
There is a caveat that the XeF6 approximate theories of chemical bonding must be viewed critically.
The values of the bond lengths and energies are average.
There are additional interactions to consider in the xenon fluorides.
The inter actions of the bond dipoles are more important in XeF21s2 than in XeF41s2.
The melting point of XeF21s2 is higher than that of XeF41s2 because of the interactions of bond dipoles.
There are other compounds in which xenon is bonding to chlorine.
The simpler compounds include XeO3 and XeOF2.
Many of the compounds have to be prepared using an indirect route.
There are possible interactions between bonds in XeF2 and XeF4 that lead to high melting points.
The other noble gases do not react directly with fluorine because the bond energy of the bond is small, and therefore the six lone formed with F is not large enough to offset the energy requirements of break pairs.
The KrF2 bond contributes only 50 kJ mol-1 to the bond energy.
We can use a similar analysis to understand why xenon doesn't react with O2 to form oxides.
The water reacts with the xenon fluorides to form various products.
XeOF4 is first hydrolyzed to xenon trioxide in the solution.
The xenon fluorides are good fluorinating agents.
In organic chemistry, xenon difluoride is used to add fluorine atoms to carbon compounds.
The by-product, Xe(g), is easily separated from the desired product by using XeF2 for this purpose.
The xenon fluorides can be used to oxidize other elements.
XeF4 oxidizes SF4 to SF6.
The oxidation state of sulfur changes from + 4 to + 6.
The name was changed to include bromine, fluorine, and iodine as well.
The group 17 in the IUPAC table was placed in Group VII of the periodic table.
Current interest in the halogens goes beyond their ability to form metallic salts.
The melting and boiling points of these elements are relatively low.
As we move down the group from the smallest and lightest member of the group, fluorine, to the largest and heaviest member, iodine, the melting and boiling points will increase.
The halogen atoms have large electron affinities and show a tendency to gain electrons.
The halogens are good oxidizers.
The elements of period 2 have a different chemistry than the rest of the group because of their small sizes and inability to expand their valence shells.
The differences between the second-row element and the members of the group are less dramatic for the halogens.
In the perhalic acids, atoms form just one bond, whereas chlorine, HXO4 and bromine form more than one bond.
Four of the seven bonds of the halogens are found in nature only as compounds, and fluorine is more reactive than the other members of the group.
It reacts to the bonds.
Oxygen, nitrogen, and the lighter noble Lewis structure for perchloric gases form compounds with even the most unreactive metals.
The acid is shown.
The chemistry of the Main- Group Elements II: Groups 18, 17, 16, 15, and Hydrogen Fluorine has a tendency to form ionic bonds with metals.
We can see this when we look at the compounds formed by the group 13 metals.
The bonding is mostly covalent because of the larger and more polarizable chloride ion.
The fluorides of the group 13 metals are all ionic lattices, whereas the chlorides of the group 13 metals are all dimers.
The ability of fluorine to stabilization other elements in very high oxidation states is an important difference between it and the other halogens.
fluorine reacts with sulfur to give SF6, with sulfur in the +6 oxidation state, whereas chlorine reacts with molten sulfur to give S2Cl2, with sulfur in the + oxidation state.
The KEEP IN MIND reactions are part of the reaction chemistry of the halogens.
For understanding the reactivity of the halogens, standard electrode potentials are helpful.
The most reactive element of the group is fluorine.
It shows the greatest tendency to gain electrons and is the easiest to reduce.
It is not surprising that fluorine is only found in combination with other elements and the F-.
Both chlorine and bromine can be found in a variety of positive oxidation states, but they are mostly found in naturally occurring compounds.
Positive oxidation states are found in natural deposits of chlorate and perchlorate.
When we summarize the reduction tendencies of main-group metals and their ion, a few Edeg values tell the story, and these values are easily incorporated into tables, such as that in Appendix D.
The oxidation-reduction chemistry of some nonmetals is richer and involves more Edeg values.
The Edeg value for reduction of the species on the left is higher than the one on the right in these diagrams.
No one was able to create a chemical reaction to extract the free element from the compounds of fluorine.
In 1886, Henri Moissan succeeded in preparing F21g2.
Edeg values cannot be added or subtracted to obtain Edeg for the desired process because it is a reduction process and not an oxidation-reduction process.
We must first convert the Edeg values to C/rGdeg values, then combine those C/rGdeg values to get a C/rGdeg value for the desired process, and finally convert this C/rGdeg value back to a value of Edeg.
The forms of chlorine identified in Figure 22-4 are oxidation and reduction.
The total equation for the desired process is the sum of the two equations above, and the C/rGdeg value is the sum of the C/rGdeg values.
The Edeg value is related to the C/rGdeg value by the expression.
The Edeg values must be added to the C/rGdeg values when adding half-equations.
The chemistry of the Main- Group Elements II: Groups 18, 17, 16, 15, and Hydrogen.
The inventor of the electric furnace, H21g2 Moissan, won the chemistry prize in 1906.
The challenge of producing fluorine by means of a chemical reaction remained.
Although chlorine can be prepared by several chemical reactions, the usual industrial method is electroly sis of NaCl.
Salt formation in the of about 70 parts per million can be obtained from inland brine sources.
The water from the Dead Dead Sea is a good source of bromine.
A good source of bromine and a number of other chemicals can be found in the Dead Sea after the high concentrations of brine solution are adjusted to pH 3.5.
A current of air or brine is used to sweep the liberated Br2 from the water.
A bromine can be concentrated by a variety of methods.
The test for the presence of Br- is based on the reaction above.
seaweed absorb and concentrate I- selec tively in the presence of Cl- and Br- Iodine can be found in small quantities from these plants.
I2 can be obtained from inland brines in the United States.
Large deposits of NaIO3 are found in Chile.
Reducing agent is required because the oxi dation state of iodine must be reduced from + 5 in IO3 to 0 in I2
The net ionic equations for the reactions are tested in a laboratory.
The halogen elements form a variety of useful compounds and are largely used to produce them.
The halogens are used to make organic compounds.
International treaties have banned the production of chlorofluorocarbons in most countries because of the damage they do to the ozone layer.
Useful as components in harsh chemical environments, fluoroinated organic compounds tend to be chemically inert, and this makes them useful.
A variety of useful compounds have ferriine as a key element.
Some of the compounds are listed in Table 22.5.
The United States is home to the major industrial chlorine, which ranks eighth in quantity among manufactured chemicals.
It has three main uses: production of chlo chlorination of organic rinated organic compounds, production of ethylene dichloride, and compounds.
brominated organic compounds are made from bromide.
Fire retardants and pesticides are some of the things these are used for.
They are used as dyes and pharmaceuticals a lot.
AgBr is the primary light-sensitive agent used in photographic film.
Iodine is less important than chlorine.
Iodine can be used as catalysts, antiseptics, germicides, and in the preparation of pharmaceuticals and photographic emulsions.
Throughout this text, we have encountered hydrogen halides.
Strong acids in water are hydrohalic acids.
There is an explanation on page 771 for why HF is a weak acid.
The application that is highlighted in the margin is the ability to etch and dis solve glass.
The reaction is similar to one between SiO2 and HF.
It must be stored in special containers because it reacts with glass.
A method discussed in Section 21-2 is used to produce hydrogen fluoride.
The glass was etched with hydrofluoric acid.
CaSO41s2 + 2 HF1g2 is a method that works for preparing HCl but not for HI.
21g2 and F21g2 are very fast and occur with explo sive violence.
The reaction with H21g2 and Cl21g2 proceeds quickly in the presence of light.
A catalyst is required when the reaction occurs more slowly.
The data shows that the standard free energies of formation are large and negative, suggesting that the reaction goes to completion.
C/fGdeg is positive for HI(g).
This suggests that HI(g) should be separated from its elements at room temperature.
In the absence of a catalyst, the dissociation of HI(g) is very slow.
HI(g) is stable at room temperature.
The reverse of the formation reaction and the negative of the dissociation reaction are listed in Table 22.6.
The relationship C/rGdeg is called -RT ln K.
Oxoacids and Oxoanions florine have the -1 oxidation state in their compounds.
The variability of oxidation states is emphasized by the oxoacids listed in Table 22.7.
H atoms are bonded to O atoms in all of these acids.
HOClO would be a more accurate representation of the other acids.
Chlorine forms a complete set of oxoacids in the oxidation states, but bromine and iodine do not.
Only a few of the oxoacids can be isolated in pure form.