The noble gases stand in stark contrast to the rest of the p-block. This group of elements does not participate in most ordinary chemistry. They interact so poorly with themselves or other elements that they evaded detection until spectroscopic evidence and the ability to liquify air finally made it possible to separate and identify these elements a little more than a century ago.
When we hear the name ‘neon’, what is the first thing that comes to mind? Probably the eye-catching glow of city lights on Broadway, the Las Vegas strip, or other big-city attractions around the world.
It is a fair association, as the first neon sign was made in 1910, only about a decade after neon’s discovery.
But neon gas in its ground state is invisible. What we see in neon lighting is the result of running a high-voltage electron discharge through the gas, promoting electrons into a range of excited states. From this, they fall back into the ground state, emitting various wavelengths that mix to form the bright glow of red-orange that is the visible emission spectrum of neon.
Neon’s Boiling Point and Liquid State
Being the next smallest of the noble gases, neon has a boiling point closer to that of helium.
Neon’s boiling point is close enough to helium’s that in certain applications where the extreme cold of helium isn’t necessary, more affordable and abundant neon can be substituted. This helps to preserve the precious helium for only applications that require it.
An interesting point to note is that neon is a liquid only over a very narrow temperature range, between 27 Kelvin and 24.5 kelvin. Stray below that narrow range, and neon becomes a solid.
In fact, all noble gases stay liquid over just a very narrow range of temperatures before transitioning on to a solid state.
Krypton is pretty similar to argon. It is another contributor to what most people call ‘neon’ lighting, where krypton contributes a blue-white set of wavelengths. It’s used in tungsten filament light bulbs. In fact, use of the more-dense krypton reduces evaporation of the tungsten, making incandescent bulbs somewhat more efficient.
In 1933, Linus Pauling, creator of the electronegativity series, predicted that krypton’s outer electrons, due to shielding from its inner layers of electrons, just might react with fluorine. Pauling had shown that fluorine is the most electronegative element on the table, and could therefore be regarded as the most reactive element
And in the 1960s, in a lab at Temple University in Pennsylvania, the irresistible force of fluorine met the immovable object of krypton, and Pauling was proven correct.
That’s when it was discovered that krypton can, indeed react under very special circumstances to form compounds like krypton difluoride.
This is because Krypton’s valence electrons in the fourth energy level are screened by more interior electrons than the lighter noble gases. This screening means that, in spite of its full octet, a krypton atom can be coaxed to share some of its valence electrons with fluorine atoms, the most electronegative element on the table.
This is the reason why tables depicting Pauling electronegativity contain values for krypton and xenon, but no values at all for their smaller cousins. This is because not even the unmatched hunger of fluorine for chemical bonds is enough to pry electrons from the clouds of helium, neon, and argon.
This article comes directly from content in the video series Understanding the Periodic Table. Watch it now, on Wondrium.
Like all of the lighter noble gases, xenon is also useful in lighting applications because of its inertness and its emission spectrum.
Xenon’s large, complex electron cloud normally results in an emission spectrum with many lines that span the visible spectrum, making it appear nearly white when it is stimulated by an electrical current. This has made xenon particularly useful in applications requiring light that simulates daylight. The development of xenon lamps in the mid 1900s set the stage for their use in photographic flashes.
Perhaps the most appealing aspect of this newer technology is that there is no chemical reaction taking place during the flash. So, a xenon lamp can be used over and over again without consuming the flash materials.
Xenon’s daylight-like emission spectrum has made it a popular medium for lighting from automobile headlamps to theater projection bulbs.
Xenon in Medicine
Xenon has also found use in medicine—as an anesthetic. Xenon’s large electron cloud can make more significant Van der Waals attractions with other atoms and molecules around it.
This makes it more soluble than its smaller noble gas cousins in a host of materials, including blood. It allows inhaled xenon to dissolve into the bloodstream and reach the brain where it inhibits neurotransmitters.
Xenon shocked everyone by participating in chemical reactions with certain electronegative elements. Xenon’s extra layer of interior electrons screens its valance shell even more than krypton. This implies that xenon can react with fluorine under quite ordinary room temperatures while activated only by sunlight. It can even react with oxygen to form oxides.
Radon was the last of the naturally occurring noble gases to be discovered. This shouldn’t be too surprising, since radon’s most stable isotope has a half-life of just about 3 days.
Radon combines the health risks of radioactivity with the mobility of a gas. This combination makes radon anything but ‘ideal’ when it accumulates in anyone’s basement. This is because radon gas has a huge atomic weight, with a density roughly eight times that of air. So if radon gas happens to form under a house, it can easily remain trapped there for long periods of time.
Needless to say, avoiding long-term radon exposure is important for our health.
Common Questions about Noble Gases
All noble gases stay liquid over just a very narrow range of temperatures before transitioning on to a solid state.
Xenon is useful in lighting applications because of its inertness and its emission spectrum.
Linus Pauling, creator of the electronegativity series, predicted that krypton’s outer electrons, due to shielding from its inner layers of electrons, just might react with fluorine.