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22 -- Part 1: Chemistry of the Main Group Elements II:
Discuss the general trends in the 22-3 Group 17: The Halogens 22-6 Hydrogen.
Explain why the melting points of the most common xenon fluorides do not follow the trend expected of nonpolar molecules.
Discuss the differences between oxygen and sulfur.
Discuss the importance of the synthesis of ammonia from nitrogen and the allotropes of phosphorus.
There are some important uses of hydrogen and the types of hydrides formed by it.
The reactant bromine is used in the synthesis of flame retardants.
We will start with a survey of the noble gases, the atoms of which have filled valence shells.
The noble nature of these gases is due to the fact that they are unreactive, though not completely so.
We will look at the halogens and then look at groups 16 and 15.
The tendency to form more than one covalent bond increases when the shell of an atom decreases.
We will consider hydrogen in this chapter.
Hydrogen is not easy to place in the periodic table.
Some chemists prefer to use a version of the periodic table that shows hydrogen separated from the rest of the table at the top of the page.
The inside front cover has hydrogen separated from the rest of the elements.
It doesn't show the actinides or the lanthanides.
The unique nature of hydrogen makes it difficult to place it in a specific group.
The concepts emphasized in this chapter are atomic, physical, and thermodynamics, bonding and structure, acid-base chemistry, and oxidation states.
A series of compounds with a common element, such as fluorine or oxygen, can be used to uncover trends in bonding.
The general formula is used to consider the fluorides of the second- and third-row elements.
Table 22.1 shows the formulas, bonding types, and phases at room temperature of the second- and third-period elements.
As we move left to right, we can see a transition from ionic bonding to covalent bonding.
BeF2 and AlF3 are network covalents.
A giant molecule is formed in a network of bonds between atoms.
The attraction among individual molecule is due to the bonds between the atoms of the molecule and the intermolecular forces, such as dipole-dipole or London dispersion forces.
The relationship between aluminum and beryllium is shown in Table 22.1.
The substance shown has the greatest number of fluorine atoms per atom of A.
The phase is at 25 degrees.
The type of bonding in the compound is reflected in the phase of the compound at room temperature.
The melting process requires the breaking of ionic bonds in the crystal lattice.
covalent bonds have to be broken in the melting process so network covalents have high melting temperatures.
The weak intermolecular forces that contribute to the attraction between the molecule are what make the gases at room temperature.
The formulas for the second-row elements are easy to understand.
The number of fluorine atoms per formula unit is the same as the number of valence electrons.
The number of fluorine atoms per formula unit is the same as the number of valence electrons required to fill the shell of carbon, nitrogen, or oxygen.
For example, a carbon atom with configuration 3He42s22p2 requires four electrons to complete its valence shell because each fluorine atom has only one unpaired valence electron.
For the fluo rides of groups 1, 2, and 13 the number of fluorine atoms per formula unit is the same as the number of valence electrons.
To get a ns2 configuration, a fluorine atom needs only one electron to complete it's valence shell, and a lithium atom needs only one electron to complete it's nucleus.
The number of valence electrons a Be or B atom must lose to attain a noble gas configuration is related to the formulas of BeF2 and BF3.
The number of fluorine atoms per formula unit is difficult to predict for the third-row elements.
As we move left to right across the third period, the oxidation state of the element bonded to fluorine increases, except in going from sulfur to chlorine.
Chlorine has an oxidation state of +7 in some of it's compounds, but it doesn't form the heptafluoride 1ClF72 with Cl in the +7 oxidation state.
The chlorine atom isn't large enough to interact with seven fluorine atoms at the same time.
This argument is supported by the fact that heptafluoride is formed by iodine, a much larger atom.
The oxides of the second and third row elements can be seen in Table 22.2.
Table 22.2 shows the acid-base properties of the oxides.
Some elements form a compound with oxygen.
The substance shown has the element in its most highly oxidation form.
The phase is at 25 degrees.
It has an oxidation state of +2.
There are diagonal relationships between Be and Al and between B and Si.
Both Be and Al form amphoteric oxides.
A transition occurs from basic oxides to acidic oxides as we move from metallic to nonmetallic elements, or as we move from the least positive elements to the most negative elements.
The oxides of the elements in groups 14 to 17 react with water to give oxoacids.
OF2 reacts with water, but not with the oxoacid.
OF2 is not an oxide.
It has an oxidation state of +2.
The oxides of Be and Al are amphoteric because they react with both acids and bases.
The oxides of the group 14 elements are affected by this.
The oxides of C and Si are acidic, the oxide of Pb is basic, and the oxide of Sn is amphoteric.
We will see these trends repeated many times as we discuss the chemistry of groups 15 through 18.
Explanations of the trends are based on differences in electronegativities and sizes of atoms.
Henry Cavendish, the discoverer of hydrogen, passed electric discharges through air to form oxides of nitrogen.
He dissolved the oxides in the water.
Cavendish wasn't able to get all the air to react by using excess oxygen.
The gas was isolated one century later by John and William.
The Congress Print and Photographs Division reasoned that there should be other members of the group because argon resembled no other element.
Ramsay began a search for other gases.
He was the first to extract helium from a mineral.
The Scottish chemist received a number of chemicals in 1904.
The last member of the group of gases was discovered in 1900.
His work on noble gases was done in 1962.
The chemistry of the Main- Group Elements II: Groups 18, 17, 16, 15, and Hydrogen inert after all.
There is a group of 18 in the periodic table.
By volume, the Occurrence Air contains 0.05% He, 01818% Ne, and 0.934% Ar.
The proportion of Kr and Xe is about 1 and 0.05 parts per million, respectively.
The atmosphere is the only place where all of these gases can be found.
Natural gas wells in the western United States region of an electric discharge that produce up to 8% He by volume are the main source of Gaseous atoms.
It costs less to emit light.
Light can be used to extract natural gas even at low levels.
Ar is the only noble gas that has escaped from the atmosphere.
The high concentration of Ar is due to the fact that it is being formed by the decay of a naturally occurring radioactive isotope.
Helium escapes from the atmosphere at a higher rate because it is 10 times less dense than Ar.
The lighter noble gases are important because of their chemical composition.
The efficiency and life of electric lightbulbs are increased when they are filled with an argon-nitrogen mixture.
Neon light is produced by electric discharge through neon-filled glass or plastic tubes.
In lasers and flashlamps, xenon and Krypton are used.
There are several unique physical properties of Helium.
For deep-sea diving, helium mixed with oxygen is used to prevent breathing during diving.
Nitrogen can be used to make powerful magnets.
Nuclear fusion research uses such magnets.
Nuclear magnetic resonance (NMR) instruments in research laboratories are one of the more familiar uses of large electromagnets.
Lighter-than-air airships are filled with helium.
The applications are highlighted in the margin.
Most of the noble gas compounds contain xenon, which is why we focus on them in this section.
KrF2 is one of the compounds that have been synthesized.
The chemistry of radon is complicated by its radioactivity, but it is expected to form compounds even more readily than xenon.
The noble gases were thought to be harmless.
The framework for the Lewis theory was provided by this apparent inertness.
It was found that compounds of xenon can be made fairly easily, and these compounds have added a lot to our knowledge of chemical bonding.
At the time, attempts were made to make oxide and fluoride compounds, but they failed.
O2 and PtF6 would combine in a mole ratio to form the compound O2PtF6.
The properties of this compound suggest it is ionic.
The first ionized energy of Xe is 1170 kJ mol-1 and the energy required to extract an electron from O2 is 1177 kJ mol-1.
The size of the Xe atom is 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- 888-609- It was thought that nature of hydrogen, helium is used in airships.
They were able to make a yellow solid with that composition.
In 2000, chemists at the noble gas compounds synthesised several more.
The conditions needed to form noble gas University of Helsinki compounds are the same as predicted by Pauling.
H4XeO6 pounds are stable at higher temperatures.
It is difficult to oxidize Xe to the positive oxidation states.
Khriachtchev, M. Pettersson, and others expect Xe compounds to be very strong oxidizing.
The significance of this large Edeg value is that XeF2 is not very stable.
All three of the xenon fluorides are volatile.
F (e) XeF6 is capped.
The synthesis of XeF2 requires an excess of Xe and the synthesis of XeF6 requires an excess of F2.
All three products tend to form at the same time.
XeF4 is the most difficult to prepare in pure form.
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