Friday, April 27, 2012

Xenon Compounds

For a long time, scientists believed that you could not make covalent compounds using Noble Gas, Group 18, elements because they are just not reactive. All these elements have a full valence shell of electrons (s2p6 for all the Noble Gases except helium, which has no p electrons) so there is no tendency to gain or lose electrons.

In 1933, Linus Pauling used his new concept of electronegativity, the power of an atom to attract electrons, to predict the possibility of making compounds using the heavier Noble Gas elements and fluorine or oxygen.

Fluorine, with an an electronegativity of 4.0 is the most electronegative element in the Periodic Table, so atoms of fluorine have the greatest power to attract electrons. Oxygen is the second most electrongetative element with an electronegativity of 3.5. So, if you want to try to make a compound with an "unreactive" non-metal like a Noble Gas, you should probably use a reactive non-metallic element like fluorine or oxygen.

Now, as you go down any Group of the Periodic Table from top to bottom, the electrons in the highest energy levels (the valence electrons) are less strongly attracted to the nucleus because they are further away from the nucleus. So, if  you want your "unreactive" Noble Gas atom to share electrons with your more reactive fluorine or oxygen atom, you will want to use a Noble Gas from the bottom of the Group, that is, xenon or radon.

So why didn't Linus Pauling predict radon compounds instead of xenon compounds then?
Radon is a radioactive element which occurs naturally as the nuclear decay product of uranium or thorium.
The real problem, as far as the production of radon compounds is concerned, is that its most stable isotope, radon-222, has a half-life of just 3.8 days, not much time for the preparation, isolation, and structural determination of your new compounds!

Xenon, on the other hand, occurs in the Earth's atmosphere in trace amounts, about 0.09 ppm (by volume) and naturally occurring xenon includes 8 stable isotopes.

In June 1962, Neil Bartlett reported that xenon reacted with platinum hexafluoride to produce xenon compounds ( a mixture of XeFPtF6 and XeFPt2F11).
In September 1962, Howard Claasen reported the production of XeF4 by reacting xenon and fluorine at high temperature.
XeF2 can actually be produced simply by exposing xenon and fluorine gases to sunlight ! 

Many xenon compounds have now been synthesized, and many of these compounds include the electronegative elements fluorine and/or oxygen.

Further Reading

Suggested Study Questions
  1. Write the  electron configuration for each of the following elements:
    • helium
    • neon
    • argon
    • fluorine
    • oxygen
  2.  Which of the Noble Gas, Group 18, elements is expected to have
    • the largest atomic number
    • the largest number of electrons in an atom
    • the largest atomic radius
    • the lowest first ionization energy
    • the lowest electronegativity
  3. Uranium-234 undergoes alpha decay to produce thorium-232. Write a nuclear decay equation for this reaction.
  4. Thorium-232 undergoes alpha decay to produce radium-228. Write a nuclear decay equation for this reaction.
  5. Radium-228 undergoes spontaneous alpha decay to produce radon-224. Write a nuclear decay equation for this reaction.
  6. If you are given 25 grams of radon-222 today, how much radon-222 will remain in 38 days time?
  7. Radon-222 emits an alpha particle to produce polonium-218. If you were give 10 grams of radon-222 19 days ago, what mass of polonium-218 has been produced since then?
  8. Write the formula for platinum hexafluoride.
  9. Give the systematic names for
    • XeF2
    • XeF4
    • XeF6
  10. Explain why no compounds of neon have so far been produced.

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