How To Use Oxygen Safely

VITAS Healthcare has provided a very comprehensive video on the precautions to take when using medical oxygen.


How To Use Oxygen Safely

Video Transcript:

“How to Safely Use Oxygen In Your Home

Now that you are using oxygen in your home, there may be some changes to your lifestyle. We want to assure you that oxygen is safe to use, but only in the right conditions.

To avoid accidents, it is important to understand the potential dangers of oxygen in your home.

Oxygen is often times not taken seriously, as many believe that oxygen is just air. Oxygen itself is colourless, odourless and tasteless. It makes up about 21% of the air we breathe.
While oxygen itself does not burn, it helps things burn faster and hotter.

The air we breathe contains about 21% oxygen. The oxygen in an oxygen cylinder such as the E cylinder we have here has purity levels of 99% oxygen or greater.
Any concentration of oxygen greater than 23.5% causes materials in your home to ignite much more easily. Oxygen lowers the temperature at which everyday household objects will start to burn. This can include your bedding, clothing, carpets, draperies and even your hair. By the way, it does not matter if the oxygen is from an oxygen cylinder, oxygen concentrator or liquid oxygen reservoir.

Oxygen from any of these sources is above 23.5% and will make most materials ignite and very low temperatures, burn quickly and twice as hot as a normal fire.

We would like to provide you with a few demonstrations of the dangers of oxygen.

Please do not attempt these demonstrations as they can be very dangerous and cause severe injury.

This is how a cigarette burns in the air we are currently breathing. Now let’s see what happens to the cigarette and how it burns when we introduce medical-grade oxygen. Notice the
cigarette is burning much more quickly, it is burning much hotter and faster than it was on room air that we breathe. You may have noticed the white hot flame that indicates a flame temperature measured in the thousands of degrees Fahrenheit.

Next, we would like to show you how an article of clothing will burn in a normal environment. Notice the article of clothing does not catch on fire very quickly. Now let’s saturate a piece of clothing with some oxygen and see what happens. Please note the piece of clothing will be above the 23.5% purity levels. We are actually using a metal pan to simulate the skin of the patient during this presentation. Please notice how much more rapidly the patient’s skin as well as clothing would catch on fire with oxygen used while smoking with oxygen in the place.

We’ve made a pipe-cleaner man that we will presume to be the patient. This patient, like many other patients, continues to smoke with oxygen in use, as he is convinced that smoking with oxygen cannot be dangerous. In fact, our patient continues to smoke in bed. When he decides he wants to have a cigarette, he takes his cannula off and lays it next to him in the bed. Our patient thinks that oxygen is just air and cannot hurt him. In this case our patient is wrong, terribly wrong. Let’s place our patient in this glass jar so we can simulate saturating the patient’s clothing with oxygen. This covering on the pipe cleaner will be his clothing. Let’s see what happens when an ember is introduced from one of his cigarettes. The fire was very impressive. However, the fire you just saw was far hotter than a normal fire. Oxygen enriched fires burn between 1500 and 3000 degrees Fahrenheit or even more depending on the materials involved. This much heat will cause serious injury which are almost always life threatening.

Sadly, this does happen in real life. Our patients have been injured and some even killed by fires caused by the misuse of oxygen. I want so badly to prevent that from happening
and it’s why I’m here right now talking to you.

Here’s a real life story from a patient who was smoking while using oxygen. He agreed to talk with us on camera to try to keep this from happening to somebody else.

Gary : “I guess I wasn’t thinking right. I was smoking with the oxygen on. And I fell asleep. And I woke up to a big blaze in my face. Ended up burning myself, different parts of my body. My nose real good. My lip a little bit. I was in a lot of pain. My wife and daughter put out the fire. Put me out. Called the ambulance. I ended up rolling around on the floor in pain. I knew that you don’t wear oxygen when you smoke and I thought I was a smart enough individual that I would never do it and get burned up.
I just thought I was smart enough and I knew that wouldn’t happen to me. But it did. Well, before, I used to just take this off and I wouldn’t even turn that off because I didn’t think
I needed to. And I would just, as long as this was away from me, I thought that was good enough. And I would just take this and throw it on the floor, make sure it was a good distance from me, and I’d smoke. Now of course I know that’s not even good enough. When I want to smoke now, first thing I do is ask somebody to turn the machine off. And once they’ve turned it off, I sit this down on the floor and then I get a cigarette and I light it. Then once I’m done with it and I put it out and it’s totally out in the ash tray, then I get this and I ask somebody to turn it back on.”

Now let’s go over some important points to remember, so you can safely use oxygen in your home.

When using or storing oxygen ,keep oxygen at least 10 feet away from open flames and other sources of heat.

Oxygen cylinders and tubing require the same consideration.

Remember, electric appliances like toaster, space heaters, hair dryers, electric blankets, and electric razors have the potential to overheat and may spark when in use.

Never smoke or allow smoking around any oxygen source and it’s tubing.

Don’t use flammable materials near oxygen, hair and aerosol sprays, paints and thinners can all pose risks. Even certain petroleum products like Vaseline and Vapor Rub can pose a danger. A spark can quickly ignite these types of products increasing the risk of fire and severe burns.

Keep oxygen tanks and tubing in well ventilated areas because oxygen tends to build up in the surrounding air and concentrate in clothes, bedding and curtains.

Never store oxygen cylinders in a closed area including closets or under a bed.

Here’s how to safely use the valve on an oxygen cylinder. All oxygen cylinder valves must be opened slowly and completely, like this. If you open it too quickly, heat is generated
and it’s called the heat of recompression. There’s enough heat to ignite anything that’s inside the valve or regulator including dust. What essentially happens is you are creating an oxygen -enriched fire inside the oxygen regulator and the valve. So remember to always open the valve slowly and completely. Additionally, never use any type of lubricant or oil
on the oxygen cylinder or regulator.

One hazard of oxygen is the vast amount of energy that is stored in a full or partially full oxygen cylinder. All oxygen cylinders should be considered sleeping giants. If a large cylinder such as an “H” cylinder or “M” cylinder as seen here were to accidently fall over, the valve could break off, and the cylinder will accelerate up to 40 miles per hour in just 0.5 of a second. These large cylinders have enough energy in them to go through two cinder block walls. All oxygen cylinders should always be secured. This can be in either a rack or a stand to prevent them from falling over. This is exactly why we tell our patients to never story oxygen cylinders in the trunk of their cars. Should you be involved in a rear end collision with oxygen cylinders in the trunk the valve could be sheared off causing the cylinder to hurt you and even those outside of your car.

Because oxygen tanks have so much energy and pressure inside, it’s important to take great caution when traveling with them in a car. Secure your oxygen cylinder in the back passenger area of your vehicle. Don’t let a cylinder roll or bang around. And never keep an oxygen in the trunk. Your vehicle should be well ventilated. Leave the window slightly cracked open to prevent oxygen and hot temperatures from building up inside.

Here’s an oxygen safety summary and a few more tips.

Always follow these precautionary measures for the safety of your family and yourself.  Never smoke while using oxygen. Warn visitors not to smoke near you when you are using oxygen.

Not following these measures could result in severe injuries, total destruction of your home, and even death. Smoking while using oxygen not only endangers your life but those who
live in the home with you and people around you.

No Smoking signs should be placed on your homes main entrance.

Firefighters, emergency personnel and visitors need to know that oxygen is being used in your home.

For oxygen concentrators only use a properly grounded wall outlet. Do not place the electrical cord or oxygen tubing under rugs, carpets, blankets, cushions or furniture.

Keep an all-purpose fire extinguisher nearby and make sure you know how to use it.

Have working smoke detectors on hand and check them monthly.

Oxygen is a great benefit for patients who need oxygen therapy. However, it should always be handled with caution.

Thank you.”


Oxygen Safety

How oxygen concentrators work

Let’s learn about OGANESSON!


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Oganesson is a (produced by people/not naturally-occurring) chemical element with the symbol Og and atomic number 118. It was first made/created in 2002 at the Joint Institute for Nuclear Research (JINR) in Dubna, near Moscow, Russia, by a combined team of Russian and American scientists. In December 2015, it was recognized as one of four new elements by the Joint Working Party of the international scientific bodies IUPAC and IUPAP. It was formally named on 28 November 2016.[16][17] The name honors the nuclear physicist Yuri Oganessian, who played a leading role in the discovery of the heaviest elements in the list of all elements. It is one of only two elements named after a person who was alive at the time of naming, the other being seaborgium, and the only element whose eponym is alive today.[18]

Oganesson has the highest atomic number and highest atomic mass of all known elements. The radioactive oganesson atom is very unstable, and since 2005, only five (possibly six) atoms of the isotope oganesson-294 have been detected.[19] Although this allowed very little experimental description of its properties and possible compounds, (related to ideas about how things work or why they happen) calculations have resulted in many (statements about possible future events), including some surprising ones. For example, although oganesson is a member of group 18 (the noble gases) – the first (produced by people/not naturally-occurring) element to be so – it may be significantly (causing reactions from other people or chemicals), unlike all the other elements of that group.[3] It was before now thought to be a gas under (usual/ commonly and regular/ healthy) conditions but is now (described a possible future event) to be a solid due to relativistic effects.[3] On the list of all elements of the elements it is a p-block element and the last one of period 7.

Other than nuclear properties, no properties of oganesson or its compounds have been measured; this is due to its very limited and expensive production[26] and the fact that it (rots/becomes ruined/gets worse) very quickly. This way only (statements about possible future events) are available.

Nuclear (firm and steady nature/lasting nature/strength) and isotopes
Main article: Isotopes of oganesson

Oganesson (row 118) is (a) little above the “island of (firm and steady nature/lasting nature/strength)” (white circle) and so its centers (of cells or atoms) are (a) little more stable than otherwise (described a possible future event).
The (firm and steady nature/lasting nature/strength) of centers (of cells or atoms) quickly decreases with the increase in atomic number after curium, element 96, whose half-life is four (many, many times more/much, much less) longer than that of any later element. All nuclides with an atomic number above 101 go through (when a radioactive substance breaks down) with half-lives shorter than 30 hours. No elements with atomic numbers above 82 (after lead) have stable isotopes.[92] This is because of the ever-increasing Coulomb fear and disgust of protons, so that the strong nuclear force cannot hold the center (of a cell or atom) together against unplanned (and sudden) fission for long. Calculations suggest that without other (making steady/making firm and strong) factors, elements with more than 104 protons should not exist.[93] However, (people who work to find information) in the 1960s suggested that the closed nuclear shells around 114 protons and 184 neutrons should undo this (quality that shows weakness because important things aren’t steady or strong), creating an island of (firm and steady nature/lasting nature/strength) in which nuclides could have half-lives reaching thousands or millions of years. While scientists have still not reached the island, the mere existence of the superheavy elements (including oganesson) confirms that this (making steady/making firm and strong) effect is real, and in general the known superheavy nuclides become (more and more as time goes on) longer-lived as they approach the (described a possible future event) location of the island.[94][95] Oganesson is radioactive and has a half-life that appears to be less than a millisecond. Anyway, this is still longer than some (described a possible future event) values,[96][97] this way giving further support to the idea of the island of (firm and steady nature/lasting nature/strength).[98]

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How oxygen concentrators work

Let’s learn about YTTERBIUM!


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Ytterbium is a chemical element with the symbol Yb and atomic number 70. It is the fourteenth and next-to-the-last element in the lanthanide series, which is the basis of the relative (firm and steady nature/lasting nature/strength) of its +2 oxidation state. However, like the other lanthanides, its most common oxidation state is +3, as in its oxide, halides, and other compounds. In water-based solution, like compounds of other late lanthanides, (able to be dissolved in something) ytterbium compounds form complexes with nine water molecules. Because of its closed-shell electron setup, its density and melting and boiling points differ significantly from those of most other lanthanides.

In 1878, the Swiss chemist Jean Charles Galissard de Marignac separated from the rare earth “erbia” another independent part, which he called “ytterbia”, for Ytterby, the village in Sweden near where he found the new part of erbium. He suspected that ytterbia was a compound of a new element that he called “ytterbium” (in total, four elements were named after the village, the others being yttrium, terbium, and erbium). In 1907, the new earth “lutecia” was separated from ytterbia, from which the element “lutecium” (now lutetium) was (pulled out or taken from something else) by Georges Urbain, Carl Auer von Welsbach, and Charles James. After some discussion, Marignac’s name “ytterbium” was kept/held. A (compared to other things) total/totally/with nothing else mixed in sample of the metal was not received/got until 1953. Now, ytterbium is mainly used as a dopant of stainless steel or active laser media, and less often as a (ray of invisible energy) source.

Natural ytterbium is a mixture of seven stable isotopes, which completely are present at concentrations of 0.3 parts per million. This element is mined in China, the United States, Brazil, and India in form of the minerals monazite, euxenite, and xenotime. The ytterbium concentration is low because it is found only among many other rare-earth elements; more than that, it is among the least plentiful. Once (pulled out or taken from something else) and prepared, ytterbium is somewhat dangerous as an eye and skin irritant. The metal is a fire and explosion danger/risk.

Physical properties
Ytterbium is a soft, bendable and (able to be flattened or drawn into wire) chemical element that displays a bright silvery shine when total/totally/with nothing else mixed in. It is a rare-earth element, and it is easily (mixed with and became part of a liquid) by the strong mineral acids. It reacts slowly with cold water and it oxidizes slowly in air.[7]

Ytterbium has three give out/set asideropes labeled by the Greek letters alpha, beta and gamma; their change temperatures are aˆ’13 °C and 795 °C,[7] although the exact change temperature depends on the pressure and stress.[8] The beta give out/set asiderope (6.966 g/cm3) exists at room temperature, and it has a face-centered cubic crystal structure. The high-temperature gamma give out/set asiderope (6.57 g/cm3) has a body-centered cubic (very clear/related to things that look like little pieces of clear glass) structure.[7] The alpha give out/set asiderope (6.903 g/cm3) has a six-sided (very clear/related to things that look like little pieces of clear glass) structure and is stable at low temperatures.[9] The beta give out/set asiderope has a metallic electrical (ability to let electricity flow) at (usual/ commonly and regular/ healthy) (related to the air outside) pressure, but it becomes an (element used to make electronic circuits) when exposed to a pressure of about 16,000 atmospheres (1.6 GPa). Its electrical resistivity increases ten times upon (press or force into a smaller space)ion to 39,000 atmospheres (3.9 GPa), but then drops to about 10% of its room-temperature resistivity at about 40,000 atm (4.0 GPa).[7][10]

In contrast with the other rare-earth metals, which usually have antiferromagnetic and/or ferromagnetic properties at low temperatures, ytterbium is paramagnetic at temperatures above 1.0 kelvin.[11] However, the alpha give out/set asiderope is diamagnetic.[8] With a melting point of 824 °C and a boiling point of 1196 °C, ytterbium has the smallest liquid range of all the metals.[7]

Opposite to most other lanthanides, which have a close-packed six-sided (something made of crossed strips of wood, metal, etc.), ytterbium makes crystals/becomes clear and real in the face-centered cubic system. Ytterbium has a density of 6.973 g/cm3, which is much lower than those of the close-by lanthanides, thulium (9.32 g/cm3) and lutetium (9.841 g/cm3). Its melting and boiling points are also much lower than those of thulium and lutetium. This is due to the closed-shell electron setup of ytterbium ([Xe] 4f14 6s2), which causes only the two 6s electrons to be available for metallic (gluing or joining together of two things) (in contrast to the other lanthanides where three electrons are available) and increases ytterbium’s metallic radius.[9]

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How oxygen concentrators work

Let’s learn about DUBNIUM!


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Dubnium is a synthetic chemical element with the symbol Db and atomic number 105. Dubnium is highly radioactive: the most stable known isotope, dubnium-268, has a half-life of about 28 hours. This greatly limits extended research on dubnium.

Dubnium does not occur naturally on Earth and is produced artificially. The Soviet Joint Institute for Nuclear Research (JINR) claimed the first discovery of the element in 1968, followed by the American Lawrence Berkeley Laboratory in 1970. Both teams proposed their names for the new element and used them without formal approval. The long-standing dispute was resolved in 1993 by an official investigation of the discovery claims by the Transfermium Working Group, formed by the International Union of Pure and Applied Chemistry and the International Union of Pure and Applied Physics, resulting in credit for the discovery being officially shared between both teams. The element was formally named dubnium in 1997 after the town of Dubna, the site of the JINR.

Theoretical research establishes dubnium as a member of group 5 in the 6d series of transition metals, placing it under vanadium, niobium, and tantalum. Dubnium should share most properties, such as its valence electron configuration and having a dominant +5 oxidation state, with the other group 5 elements, with a few anomalies due to relativistic effects. A limited investigation of dubnium chemistry has confirmed this. Solution chemistry experiments have revealed that dubnium often behaves more like niobium rather than tantalum, breaking periodic trends.

Predicted properties
According to the periodic law, dubnium should belong to group 5, with vanadium, niobium, and tantalum. Several studies have investigated the properties of element 105 and found that they generally agreed with the predictions of the periodic law. Significant deviations may nevertheless occur, due to relativistic effects,[m] which dramatically change physical properties on both atomic and macroscopic scales. These properties have remained challenging to measure for several reasons: the difficulties of production of superheavy atoms, the low rates of production, which only allows for microscopic scales, requirements for a radiochemistry laboratory to test the atoms, short half-lives of those atoms, and the presence of many unwanted activities apart from those of synthesis of superheavy atoms. So far, studies have only been performed on single atoms.[3]

Atomic and physical

Relativistic (solid line) and nonrelativistic (dashed line) radial distribution of the 7s valence electrons in dubnium.
A direct relativistic effect is that as the atomic numbers of elements increase, the innermost electrons begin to revolve faster around the nucleus as a result of an increase of electromagnetic attraction between an electron and a nucleus. Similar effects have been found for the outermost s orbitals (and p1/2 ones, though in dubnium they are not occupied): for example, the 7s orbital contracts by 25% in size and is stabilized by 2.6 eV.[3]

A more indirect effect is that the contracted s and p1/2 orbitals shield the charge of the nucleus more effectively, leaving less for the outer d and f electrons, which therefore move in larger orbitals. Dubnium is greatly affected by this: unlike the previous group 5 members, its 7s electrons are slightly more difficult to extract than its 6d electrons.[3]

Relativistic stabilization of the ns orbitals, the destabilization of the (n-1)d orbitals and their spin–orbit splitting for the group 5 elements.
Another effect is the spin–orbit interaction, particularly spin–orbit splitting, which splits the 6d subshell—the azimuthal quantum number ℓ of a d shell is 2—into two subshells, with four of the ten orbitals having their ℓ lowered to 3/2 and six raised to 5/2. All ten energy levels are raised; four of them are lower than the other six. (The three 6d electrons normally occupy the lowest energy levels, 6d3/2.)[3]

A singly ionized atom of dubnium (Db+) should lose a 6d electron compared to a neutral atom; the doubly (Db2+) or triply (Db3+) ionized atoms of dubnium should eliminate 7s electrons, unlike its lighter homologs. Despite the changes, dubnium is still expected to have five valence electrons; 7p energy levels have not been shown to influence dubnium and its properties. As the 6d orbitals of dubnium are more destabilized than the 5d ones of tantalum, and Db3+ is expected to have two 6d, rather than 7s, electrons remaining, the resulting +3 oxidation state is expected to be unstable and even rarer than that of tantalum. The ionization potential of dubnium in its maximum +5 oxidation state should be slightly lower than that of tantalum and the ionic radius of dubnium should increase compared to tantalum; this has a significant effect on dubnium’s chemistry.[3]

Atoms of dubnium in the solid state should arrange themselves in a body-centered cubic configuration, like the previous group 5 elements.[4] The predicted density of dubnium is 21.6 g/cm3.[5]

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How oxygen concentrators work

Let’s learn about TERBIUM!


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Terbium is a chemical element with the symbol Tb and atomic number 65. It is a silvery-white, rare earth metal that is bendable, (able to be flattened or drawn into wire), and soft enough to be cut with a knife. The ninth member of the lanthanide series, terbium is a fairly electropositive metal that reacts with water, changing (and getting better) hydrogen gas. Terbium is never found in nature as a free element, but it is contained in many minerals, including cerite, gadolinite, monazite, xenotime, and euxenite.

Swedish chemist Carl Gustaf Mosander discovered terbium as a chemical element in 1843. He detected it as a (dirt, dust, etc.) in yttrium oxide, Y2O3. Yttrium and terbium, as well as erbium and ytterbium, are named after the village of Ytterby in Sweden. Terbium was not (separated far from others) in (completely/complete, with nothing else mixed in) form until the coming of ion exchange ways of doing things.

Terbium is used to dope (silvery metal/important nutrient) fluoride, (silvery metal/important nutrient) tungstate and strontium molybdate, materials that are used in solid-state devices, and as a crystal (make steady/make firm and strong)r of (devices that make electricity from hydrogen) which operate at high/higher temperatures. As a part of Terfenol-D (a mix/mixture (of metals) that expands and contracts when exposed to magnetic fields more than any other mix/mixture (of metals)), terbium is of use in mechanical pushing-pulling devices, in naval sonar systems and in sensors.

Most of the world’s terbium supply is used in green phosphors. Terbium oxide is in fluorescent lamps and television and monitor cathode ray tubes (CRTs). Terbium green phosphors are combined with divalent europium blue phosphors and trivalent europium red phosphors to provide trichromatic lighting technology, a high-(wasting very little while working or producing something) white light used for standard lighting up/education in indoor lighting.

Physical properties
Terbium is a silvery-white rare earth metal that is bendable, (able to be flattened or drawn into wire) and soft enough to be cut with a knife. It is staying steady in air compared to the earlier, more (causing reactions from other people or chemicals) lanthanides in the first half of the lanthanide series. Terbium exists in two crystal give out/set asideropes with a change temperature of 1289 °C between them. The 65 electrons of a terbium atom are arranged in the electron setup [Xe]4f96s2; (usually/ in a common and regular way), only three electrons can be removed before the nuclear charge becomes too great to allow further ionization, but in the case of terbium, the (firm and steady nature/lasting nature/strength) of the half-filled [Xe]4f7 setup allows further ionization of a fourth electron in the presence of very strong oxidizing agents such as fluorine gas.

The terbium(III) cation is brilliantly fluorescent, in a bright lemon-yellow color that is the result of a strong green emission line in combination with other lines in the orange and red. The yttrofluorite variety of the mineral fluorite owes its creamy-yellow fluorescence in part to terbium. Terbium easily oxidizes, and is therefore used in its elemental form specifically for research. Single terbium atoms have been (separated far from others) by inserting them into fullerene molecules.

Terbium has a simple ferromagnetic ordering at temperatures below 219 K. Above 219 K, it turns into a helical antiferromagnetic state in which all of the atomic moments in a particular basal plane layer are parallel, and oriented at a fixed angle to the moments of (next to) layers. This unusual antiferromagnetism changes into a (not working right/not acting right) paramagnetic state at 230 K.

Chemical properties
Terbium metal is an electropositive element and oxidizes in the presence of most acids (such as sulfuric acid), all of the halogens, and even water.

2 Tb (s) + 3 H2SO4 a†’ 2 Tb3+ + 3 SO2aˆ’
4 + 3 H2a†’
2 Tb + 3 X2 a†’ 2 TbX3 (X = F, Cl, Br, I)
2 Tb (s) + 6 H2O a†’ 2 Tb(OH)3 + 3 H2a†’
Terbium also oxidizes easily in air to form a mixed terbium(III,IV) oxide:[8]

8 Tb + 7 O2 a†’ 2 Tb4O7
The most common oxidation state of terbium is +3 (trivalent), such as TbCl
3. In the solid state, tetravalent terbium is also known, in compounds such as TbO2 and TbF4. In solution, terbium usually forms trivalent (group of similar living things), but can be oxidized to the tetravalent state with ozone in highly basic water-based conditions.

The coordination and organometallic chemistry of terbium is just like other lanthanides. In water-based conditions, terbium can be coordinated by nine water molecules, which are arranged in a tricapped trigonal prismatic molecular geometry. Complexes of terbium with lower coordination number are also known, usually with (taking up a lot of space for its weight) ligands like bis(trimethyl-silylamide), which forms the three-coordinate Tb[N(SiMe3)2]3 complex.

Most coordination and organometallic complexes contain terbium in the trivalent oxidation state. Divalent (Tb2+) complexes are also known, usually with (taking up a lot of space for its weight) cyclopentadienyl-type ligands. A few coordination compounds containing terbium in its tetravalent state are also known.


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Let’s learn about XENON!


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Xenon is a chemical element with the symbol Xe and atomic number 54. It is a colorless, dense, odorless noble gas found in Earth’s atmosphere in trace amounts. Although generally unreactive, xenon can undergo a few chemical reactions such as the formation of xenon hexafluoroplatinate, the first noble gas compound to be synthesized.

Xenon is used in flash lamps and arc lamps, and as a general anesthetic. The first excimer laser design used a xenon dimer molecule (Xe2) as the lasing medium, and the earliest laser designs used xenon flash lamps as pumps. Xenon is used to search for hypothetical weakly interacting massive particles[20] and as the propellant for ion thrusters in spacecraft.

Naturally occurring xenon consists of seven stable isotopes and two long-lived radioactive isotopes. More than 40 unstable xenon isotopes undergo radioactive decay, and the isotope ratios of xenon are an important tool for studying the early history of the Solar System. Radioactive xenon-135 is produced by beta decay from iodine-135 (a product of nuclear fission), and is the most significant (and unwanted) neutron absorber in nuclear reactors.

Xenon was discovered in England by the Scottish chemist William Ramsay and English chemist Morris Travers in September 1898, shortly after their discovery of the elements krypton and neon. They found xenon in the residue left over from evaporating components of liquid air. Ramsay suggested the name xenon for this gas from the Greek word ξένον xénon, neuter singular form of ξένος xénos, meaning ‘foreign(er)’, ‘strange(r)’, or ‘guest’. In 1902, Ramsay estimated the proportion of xenon in the Earth’s atmosphere to be one part in 20 million.

During the 1930s, American engineer Harold Edgerton began exploring strobe light technology for high speed photography. This led him to the invention of the xenon flash lamp in which light is generated by passing brief electric current through a tube filled with xenon gas. In 1934, Edgerton was able to generate flashes as brief as one microsecond with this method.

In 1939, American physician Albert R. Behnke Jr. began exploring the causes of “drunkenness” in deep-sea divers. He tested the effects of varying the breathing mixtures on his subjects, and discovered that this caused the divers to perceive a change in depth. From his results, he deduced that xenon gas could serve as an anesthetic. Although Russian toxicologist Nikolay V. Lazarev apparently studied xenon anesthesia in 1941, the first published report confirming xenon anesthesia was in 1946 by American medical researcher John H. Lawrence, who experimented on mice. Xenon was first used as a surgical anesthetic in 1951 by American anesthesiologist Stuart C. Cullen, who successfully used it with two patients.

Xenon and the other noble gases were for a long time considered to be completely chemically inert and not able to form compounds. However, while teaching at the University of British Columbia, Neil Bartlett discovered that the gas platinum hexafluoride (PtF6) was a powerful oxidizing agent that could oxidize oxygen gas (O2) to form dioxygenyl hexafluoroplatinate (O+2[PtF 6]−). Since O2(1165 kJ/mol) and xenon (1170 kJ/mol) have almost the same first ionization potential, Bartlett realized that platinum hexafluoride might also be able to oxidize xenon. On March 23, 1962, he mixed the two gases and produced the first known compound of a noble gas, xenon hexafluoroplatinate.

Bartlett thought its composition to be Xe+[PtF6]−, but later work revealed that it was probably a mixture of various xenon-containing salts. Since then, many other xenon compounds have been discovered, in addition to some compounds of the noble gases argon, krypton, and radon, including argon fluorohydride (HArF), krypton difluoride (KrF2), and radon fluoride. By 1971, more than 80 xenon compounds were known.

In November 1989, IBM scientists demonstrated a technology capable of manipulating individual atoms. The program, called IBM in atoms, used a scanning tunneling microscope to arrange 35 individual xenon atoms on a substrate of chilled crystal of nickel to spell out the three letter company initialism. It was the first time atoms had been precisely positioned on a flat surface.

Xenon has atomic number 54; that is, its nucleus contains 54 protons. At standard temperature and pressure, pure xenon gas has a density of 5.761 kg/m3, about 4.5 times the density of the Earth’s atmosphere at sea level, 1.217 kg/m3. As a liquid, xenon has a density of up to 3.100 g/mL, with the density maximum occurring at the triple point. Liquid xenon has a high polarizability due to its large atomic volume, and thus is an excellent solvent. It can dissolve hydrocarbons, biological molecules, and even water. Under the same conditions, the density of solid xenon, 3.640 g/cm3, is greater than the average density of granite, 2.75 g/cm3. Under gigapascals of pressure, xenon forms a metallic phase.


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Let’s learn about TELLURIUM!


Native Tellurium


Tellurium is a chemical element with the symbol Te and atomic number 52. It is a brittle, mildly toxic, rare, silver-white metalloid. Tellurium is chemically related to selenium and sulfur, all three of which are chalcogens. It is occasionally found in native form as elemental crystals. Tellurium is far more common in the Universe as a whole than on Earth. Its extreme rarity in the Earth’s crust, comparable to that of platinum, is due partly to its formation of a volatile hydride that caused tellurium to be lost to space as a gas during the hot nebular formation of Earth, and partly to tellurium’s low affinity for oxygen, which causes it to bind preferentially to other chalcophiles in dense minerals that sink into the core.

Tellurium-bearing compounds were first discovered in 1782 in a gold mine in Kleinschlatten, Transylvania (now Zlatna, Romania) by Austrian mineralogist Franz-Joseph Müller von Reichenstein, although it was Martin Heinrich Klaproth who named the new element in 1798 after the Latin word for “earth”, tellus. Gold telluride minerals are the most notable natural gold compounds. However, they are not a commercially significant source of tellurium itself, which is normally extracted as a by-product of copper and lead production.

Commercially, the primary use of tellurium is copper (tellurium copper) and steel alloys, where it improves machinability. Applications in CdTe solar panels and cadmium telluride semiconductors also consume a considerable portion of tellurium production. Tellurium is considered a technology-critical element.

Tellurium has no biological function, although fungi can use it in place of sulfur and selenium in amino acids such as tellurocysteine and telluromethionine. In humans, tellurium is partly metabolized into dimethyl telluride, (CH3)2Te, a gas with a garlic-like odor exhaled in the breath of victims of tellurium exposure or poisoning.

Tellurium has two allotropes, crystalline and amorphous. When crystalline, tellurium is silvery-white with a metallic luster. It is a brittle and easily pulverized metalloid. Amorphous tellurium is a black-brown powder prepared by precipitating it from a solution of tellurous acid or telluric acid (Te(OH)6). Tellurium is a semiconductor that shows a greater electrical conductivity in certain directions depending on atomic alignment; the conductivity increases slightly when exposed to light (photoconductivity). When molten, tellurium is corrosive to copper, iron, and stainless steel. Of the chalcogens (oxygen-family elements), tellurium has the highest melting and boiling points, at 722.66 K (841.12 °F) and 1,261 K (1,810 °F), respectively.

Tellurium adopts a polymeric structure consisting of zig-zag chains of Te atoms. This gray material resists oxidation by air and is not volatile.

Naturally occurring tellurium has eight isotopes. Six of those isotopes, 120Te, 122Te, 123Te, 124Te, 125Te, and 126Te, are stable. The other two, 128Te and 130Te, have been found to be slightly radioactive, with extremely long half-lives, including 2.2 × 1024 years for 128Te. This is the longest known half-life among all radionuclides and is about 160 trillion (1012) times the age of the known universe. Stable isotopes comprise only 33.2% of naturally occurring tellurium.

A further 31 artificial radioisotopes of tellurium are known, with atomic masses ranging from 104 to 142 and with half-lives of 19 days or less. Also, 17 nuclear isomers are known, with half-lives up to 154 days. With the exception of beryllium-8 and beta-delayed alpha emission branches in some lighter nuclides, tellurium (104Te to 109Te) is the lightest element with isotopes known to undergo alpha decay.

The atomic mass of tellurium (127.60 g·mol−1) exceeds that of iodine (126.90 g·mol−1), the next element in the periodic table.




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Let’s learn about POLONIUM!




Polonium is a chemical element with the symbol Po and atomic number 84. Polonium is a chalcogen. A rare and highly radioactive metal with no stable isotopes, polonium is chemically similar to selenium and tellurium, though its metallic character resembles that of its horizontal neighbors in the periodic table: thallium, lead, and bismuth. Due to the short half-life of all its isotopes, its natural occurrence is limited to tiny traces of the fleeting polonium-210 (with a half-life of 138 days) in uranium ores, as it is the penultimate daughter of natural uranium-238. Though slightly longer-lived isotopes exist, they are much more difficult to produce. Today, polonium is usually produced in milligram quantities by the neutron irradiation of bismuth. Due to its intense radioactivity, which results in the radiolysis of chemical bonds and radioactive self-heating, its chemistry has mostly been investigated on the trace scale only.

Polonium was discovered in July 1898 by Marie and Pierre Curie, when it was extracted from the uranium ore pitchblende and identified solely by its strong radioactivity: it was the first element to be so discovered. Polonium was named after Marie Curie’s homeland of Poland. Polonium has few applications, and those are related to its radioactivity: heaters in space probes, antistatic devices, sources of neutrons and alpha particles, and poison. It is extremely dangerous to humans.

210Po is an alpha emitter that has a half-life of 138.4 days; it decays directly to its stable daughter isotope, 206Pb. A milligram (5 curies) of 210Po emits about as many alpha particles per second as 5 grams of 226Ra. A few curies (1 curie equals 37 gigabecquerels, 1 Ci = 37 GBq) of 210Po emit a blue glow which is caused by ionisation of the surrounding air.

About one in 100,000 alpha emissions causes an excitation in the nucleus which then results in the emission of a gamma ray with a maximum energy of 803 keV.

Polonium is a radioactive element that exists in two metallic allotropes. The alpha form is the only known example of a simple cubic crystal structure in a single atom basis at STP, with an edge length of 335.2 picometers; the beta form is rhombohedral.] The structure of polonium has been characterized by X-ray diffraction and electron diffraction.

210Po (in common with 238Pu[citation needed]) has the ability to become airborne with ease: if a sample is heated in air to 55 °C (131 °F), 50% of it is vaporized in 45 hours to form diatomic Po2 molecules, even though the melting point of polonium is 254 °C (489 °F) and its boiling point is 962 °C (1,764 °F). More than one hypothesis exists for how polonium does this; one suggestion is that small clusters of polonium atoms are spalled off by the alpha decay.

The chemistry of polonium is similar to that of tellurium, although it also shows some similarities to its neighbor bismuth due to its metallic character. Polonium dissolves readily in dilute acids but is only slightly soluble in alkalis. Polonium solutions are first colored in pink by the Po2+ ions, but then rapidly become yellow because alpha radiation from polonium ionizes the solvent and converts Po2+ into Po4+. As polonium also emits alpha-particles after disintegration so this process is accompanied by bubbling and emission of heat and light by glassware due to the absorbed alpha particles; as a result, polonium solutions are volatile and will evaporate within days unless sealed. At pH about 1, polonium ions are readily hydrolyzed and complexed by acids such as oxalic acid, citric acid, and tartaric acid.

Polonium has no common compounds, and almost all of its compounds are synthetically created; more than 50 of those are known. The most stable class of polonium compounds are polonides, which are prepared by direct reaction of two elements. Na2Po has the antifluorite structure, the polonides of Ca, Ba, Hg, Pb and lanthanides form a NaCl lattice, BePo and CdPo have the wurtzite and MgPo the nickel arsenide structure. Most polonides decompose upon heating to about 600 °C, except for HgPo that decomposes at ~300 °C and the lanthanide polonides, which do not decompose but melt at temperatures above 1000 °C. For example, PrPo melts at 1250 °C and TmPo at 2200 °C. PbPo is one of the very few naturally occurring polonium compounds, as polonium alpha decays to form lead.

Polonium hydride (PoH2) is a volatile liquid at room temperature prone to dissociation; it is thermally unstable. Water is the only other known hydrogen chalcogenide which is a liquid at room temperature; however, this is due to hydrogen bonding. The three oxides, PoO, PoO2 and PoO3, are the products of oxidation of polonium.

Halides of the structure PoX2, PoX4 and PoF6 are known. They are soluble in the corresponding hydrogen halides, i.e., PoClX in HCl, PoBrX in HBr and PoI4 in HI. Polonium dihalides are formed by direct reaction of the elements or by reduction of PoCl4 with SO2 and with PoBr4 with H2S at room temperature. Tetrahalides can be obtained by reacting polonium dioxide with HCl, HBr or HI.

Other polonium compounds include potassium polonite as a polonite, polonate, acetate, bromate, carbonate, citrate, chromate, cyanide, formate, (II) and (IV) hydroxides, nitrate, selenate, selenite, monosulfide, sulfate, disulfate and sulfite.

A limited organopolonium chemistry is known, mostly restricted to dialkyl and diaryl polonides (R2Po), triarylpolonium halides (Ar3PoX), and diarylpolonium dihalides (Ar2PoX2). Polonium also forms soluble compounds with some chelating agents, such as 2,3-butanediol and thiourea.



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Let’s learn about BISMUTH!

Bismuth Crystal | Rainbow Bismuth - Rose and Rabbit Designs

Bismuth is a chemical element with the symbol Bi and atomic number 83. It is a pentavalent post-transition metal and one of the pnictogens with chemical properties resembling its lighter group 15 siblings arsenic and antimony. Elemental bismuth may occur naturally, although its sulfide and oxide form important commercial ores. The free element is 86% as dense as lead. It is a brittle metal with a silvery-white color when freshly produced, but surface oxidation can give it an iridescent tinge in numerous colours. Bismuth is the most naturally diamagnetic element and has one of the lowest values of thermal conductivity among metals.

Bismuth was long considered the element with the highest atomic mass that is stable, but in 2003 it was discovered to be extremely weakly radioactive: its only primordial isotope, bismuth-209, decays via alpha decay with a half-life more than a billion times the estimated age of the universe. Because of its tremendously long half-life, bismuth may still be considered stable for almost all purposes.

Bismuth metal has been known since ancient times, although it was often confused with lead and tin, which share some physical properties. The etymology is uncertain, but the word may come from the German words weiße Masse or Wismuth (“white mass”), translated in the mid-sixteenth century to New Latin bisemutum or bisemutium.

Bismuth compounds account for about half the production of bismuth. They are used in cosmetics; pigments; and a few pharmaceuticals, notably bismuth subsalicylate, used to treat diarrhea. Bismuth’s unusual propensity to expand as it solidifies is responsible for some of its uses, such as in the casting of printing type. Bismuth has unusually low toxicity for a heavy metal. As the toxicity of lead has become more apparent in recent years, there is an increasing use of bismuth alloys (presently about a third of bismuth production) as a replacement for lead.

Bismuth metal has been known since ancient times; it was one of the first 10 metals to have been discovered. The name bismuth dates from around the 1660s and is of uncertain etymology; it possibly comes from obsolete German Bismuth, Wismut, Wissmuth (early 16th century), perhaps related to Old High German hwiz (“white”). The New Latin bisemutium (due to Georgius Agricola, who Latinized many German mining and technical words) is from the German Wismuth, perhaps from weiße Masse, “white mass”.

The element was confused in early times with tin and lead because of its resemblance to those elements. Because bismuth has been known since ancient times, no one person is credited with its discovery. Agricola (1546) states that bismuth is a distinct metal in a family of metals including tin and lead. This was based on observation of the metals and their physical properties.

Miners in the age of alchemy also gave bismuth the name tectum argenti, or “silver being made,” in the sense of silver still in the process of being formed within the Earth.

Bismuth was also known to the Incas and used (along with the usual copper and tin) in a special bronze alloy for knives.

Alchemical symbol used by Torbern Bergman (1775).  Beginning with Johann Heinrich Pott in 1738, Carl Wilhelm Scheele, and Torbern Olof Bergman, the distinctness of lead and bismuth became clear, and Claude François Geoffroy demonstrated in 1753 that this metal is distinct from lead and tin.

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