P__ by Pramod Kothari

P-block elements

Lulu Press, Raleigh, N.C. USA

Dr. Pramod Kothari Assistant Professor, Department Of Chemistry Government Post Graduate College, Berinag, District â&#x20AC;&#x201C; Pithoragarh Uttarakhand (India)

Copyright ÂŠ Creative Commons Attribution-Share Alike 3.0 //creativecommons.org/licenses/by-sa/3.0/ Disclaimer All the material contained in this book is provided for educational and informational purposes only. No responsibility can be taken for any results or outcomes resulting from the use of this material. While every attempt has been made to provide information that is both accurate and effective, the author does not assume any responsibility for the accuracy or use/misuse of this information.

Dr. Pramod Kothari / Alkali Metals, ISBN/EAN13: 978-1-304-87546-4

Preface The p-block of the periodic table of the elements consists of the last six groups except helium (which is located in the s-block). In the elemental form of the p-block elements, the highest energy electron occupies a p-orbital. The p-block contains all of the nonmetals (except for hydrogen and helium, which are in the s-block) and semimetals, as well as the poor metals. The groups of the p-block are: 

13 (IIIB, IIIA): Boron group



14 (IVB, IVA): Carbon group



15 (VB, VA): Nitrogen group (or pnictogens)



16 (VIB, VIA): Chalcogens



17 (VIIB, VIIA): Halogens



18 (0, VIIIA): Noble gases (excluding helium)

Many of the p-block elements have been known since antiquity, and all naturally occurring pblock elements with the exception of astatine were discovered before 1900. Astatine was finally discovered in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio Segrè at the University of California, Berkeley. The remaining p-block elements are hypothesized, based on periodic trends, to be elements 113–118, although it is currently unknown if they are actually p-block elements.

Dr. Pramod Kothari Assistant Professor, Department Of Chemistry Government Post Graduate College, Berinag, District – Pithoragarh Uttarakhand (India)

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Table of Contents P-block .......................................................................................................... 1 Boron Group .................................................................................................. 4 Carbon Group Section ................................................................................. 18 Pnictogen .................................................................................................... 28 Chalcogen .................................................................................................... 37 Halogen. ...................................................................................................... 53 Noble gas ..................................................................................................... 67

Dr. Pramod Kothari / Alkali Metals, ISBN/EAN13: 978-1-304-87546-4

Chapter 1: P-block p-block in the periodic table

Boron

Carbon

Nitrogen

Oxygen

Fluorine

Neon

Silicon

Phosphorus

Sulfur

Chlorine

Argon

Arsenic

Selenium

Bromine

Krypton

Group →

↓ Period

3 Aluminium

4 Gallium Germanium

Indium

Tin

Antimony

Tellurium

Iodine

Xenon

Lead

Bismuth

Polonium

Astatine

Radon

6 Thallium 81

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7 Ununtrium Flerovium Ununpentium Livermorium Ununseptium Ununoctium 113

114

115

116

117

118

Uut

Uup

Uus

Uuo

The p-block is one of two blocks in the periodic table to contain nonmetals (although the sblock only contains two nonmetals, hydrogen and helium). As such, it has some of the most diverse properties of any region in the periodic table. The metals in this region of the periodic table are, in general, softer and have lower melting points than transition metals. The p-block is the only region of the periodic table to contain metalloids. In general, the farther one goes to the right, and the farther one goes up in the p-block, the less metallic the elements get; the metalloids form a diagonal line from the upper left to the lower right of the p-block. All elements in the p-block have their outermost electron in a p-subshell. Trend from metal to nonmetal through metalloid in the p-block 13

Group â&#x2020;&#x2019; â&#x2020;&#x201C; Period

5 B

6 C

10 Ne

8 O

3 15 14 13

17 16 S

18 Ar

33 34

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85 At

81 Tl

Legend for the periodic table

black=Solid

green=Liquid

red=Gas

Color of the atomic number shows state of matter (at 0 Â°C and 1 atm)

Primordial From decay Synthetic

Border shows natural occurrence of the element

Background color shows subcategory in the metalâ&#x20AC;&#x201C;nonmetal range: Metal

Metalloid Nonmetal

Alkali

Alkaline

Lan-

metal

earth metal thanide

Actinide Transition metal

Poor

Polyatomic

Diatomic

Noble

metal

nonmetal

gas

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Chapter 2: Boron group

Boron group (group 13)

group 12 ←

→ carbon group

IUPAC group number

Name by element

boron group

Trivial name

triels, icosagens

CAS group number (US)

IIIA

old IUPAC number (European)IIIB

↓ Period

Boron (B)

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5 Metalloid

Aluminium (Al) 13 Poor metal

Gallium (Ga) 31 Poor metal

Indium (In) 49 Poor metal

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Thallium (Tl) 81 Poor metal

Ununtrium (Uut) 113 unknown chemical properties

Legend primordial element synthetic element Atomic number color: black=solid

The 5 stable elements of the boron group

The boron group are the chemical elements in group 13 of the periodic table, comprising boron (B), aluminium (Al), gallium (Ga), indium (In), thallium (Tl), and ununtrium (Uut). The elements in the boron group are characterized by having three electrons in their outer energy levels (valence layers). These elements have also been referred to as icosagens and triels. Boron is classified as a metalloid while the rest, with the possible exception of ununtrium, are considered poor metals. Ununtrium has not yet been confirmed to be a poor metal and, due to relativistic effects, might not turn out to be one. Boron occurs sparsely, probably because bombardment by the subatomic particles produced from natural radioactivity disrupts its

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nuclei. Aluminium occurs widely on earth, and indeed is the third most abundant element in the Earth's crust (8.3%). Gallium is found in the earth with an abundance of 13 ppm. Indium is the 61st most abundant element in the earth's crust, and thallium is found in moderate amounts throughout the planet. Ununtrium is never found in nature and therefore is termed a synthetic element. Several group 13 elements have biological roles in the ecosystem. Boron is a trace element in humans and is essential for some plants. Lack of boron can lead to stunted plant growth, while an excess can also cause harm by inhibiting growth. Aluminium has neither a biological role nor significant toxicity and is considered safe. Indium and gallium can stimulate metabolism; gallium is credited with the ability to bind itself to iron proteins. Thallium is highly toxic, interfering with the function of numerous vital enzymes, and has seen use as a pesticide. Characteristics Like other groups, the members of this family show patterns in electron configuration, especially in the outermost shells, resulting in trends in chemical behavior: Z

Element No. of electrons per shell

boron

2, 3

13 aluminium 2, 8, 3 31 gallium

2, 8, 18, 3

49 indium

2, 8, 18, 18, 3

81 thallium

2, 8, 18, 32, 18, 3

113 ununtrium 2, 8, 18, 32, 32, 18, 3 The boron group is notable for trends in the electron configuration, as shown above, and in some of its elements' characteristics. Boron differs from the other group members in its hardness, refractivity and reluctance to participate in metallic bonding. An example of a trend in reactivity is boron's tendency to form reactive compounds with hydrogen. Chemical reactivity

Hydrides Most of the elements in the boron group show increasing reactivity as the elements get heavier in atomic mass and higher in atomic number. Boron, the first element in the group, is generally unreactive with many elements except at high temperatures, although it is capable of forming many compounds with hydrogen, sometimes called boranes. The simplest borane is diborane, or B2H6. Another example is B10H14.

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The next group-13 elements, aluminium and gallium, form fewer stable hydrides, although both AlH3 and GaH3 exist. Indium, the next element in the group, is not known to form many hydrides, except in complex compounds such as the phosphine complex H 3InP(Cy)3. No stable compound of thallium and hydrogen has been synthesized in any laboratory. Some common chemical compounds of the boron group Element Oxides Boron

Hydrides

Fluorides Chlorides Sulfides

(β/g/α)B2O3

B2H6

BF3

B2O

B10H14

BF−

BCl3

B2S3

AlCl3

(α/β/γ) Al2S3

GaCl3

GaS

4 B6O

BH3

B2F4

B5H9

B6H12 B4H10 B6H2−6 B12H2−12 B20H26 Aluminium (γ/δ/η/θ/σ)Al2O3 (α/α`/β/δ/ε/θ/γ) AlH3 AlF3 Al2O

Al2H6

AlO

AlH4 AlH− 4

Gallium

(α/β/δ/γ/ε) Ga2O3 Ga2H6

GaF3

GaH4

GaCl2

GaH3

Ga2Cl4

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Ga2Cl6 −4

GaCl

−7

Ga2Cl Indium

In2O3

InH3

InF3

InCl3

Tl2O3

TlH3

TlF

TlCl

Tl2O

TlH

TlF3

TlCl3

(α/β/γ) In2S3

In2O Thallium

TlO2

TlF

Tl4O3

TlF

Ununtrium Uut2O Uut2O3

3−4

2−3

TlCl2 Tl2Cl3

UutH

UutF

UutCl

UutH3

UutF3

UutCl3

UutF5 UutF

−6

Oxides All of the boron-group elements are known to form a trivalent oxide, with two atoms of the element bonded covalently with three atoms of oxygen. These elements show a trend of increasing pH (from acidic to basic).Boron oxide (B 2O3) is slightly acidic, aluminium and gallium oxide (Al2O3 and Ga2O3 respectively) are amphoteric, indium(III) oxide (In2O3) is nearly amphoteric, and thallium(III) oxide (Tl2O3) is a Lewis base because it dissolves in acids to form salts. Each of these compounds are stable, but thallium oxide decomposes at temperatures higher than 875 °C.

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A powdered sample of boron trioxide (B2O3), one of the oxides of boron

Halides The elements in group 13 are also capable of forming stable compounds with the halogens, usually with the formula MX3 (where M is a boron-group element and X is a halogen.) The only exception to this is thallium(III) iodide.Fluorine, the first halogen, is able to form stable compounds with every element that has been tested (except neon and helium), and the boron group is no exception. It is even hypothesized that ununtrium could form a compound with fluorine, UutF3, before spontaneously decaying due to ununtrium's radioactivity. Chlorine also forms stable compounds with all of the elements in the boron group, including thallium, and is hypothesized to react with ununtrium. All of the elements will react with bromine under the right conditions, as with the other halogens but less vigorously than either chlorine or fluorine. Iodine will react with all natural elements in the periodic table except for the noble gases, and is notable for its explosive reaction with aluminium to form 2AlI 3.Astatine, the heaviest halogen, has only formed a few compounds, due to its radioactivity and short halflife, and no reports of a compound with an At–B, –Al, –Ga, –In, –Tl, or –Uut bond have been seen, although scientists think that it should form salts with metals. Physical properties It has been noticed that the elements in the boron group have similar physical properties, although most of boron's are exceptional. For example, all of the elements in the boron group, except for boron itself, are soft. Moreover, all of the other elements in group 13 are relatively reactive at moderate temperatures, while boron's reactivity only becomes comparable at very high temperatures. One characteristic that all do have in common is having three electrons in their valence shells. Boron, being a metalloid, is a thermal and electrical insulator at room temperature, but a good conductor of heat and electricity at high temperatures. Unlike boron, the metals in the group are good conductors under normal conditions. This is in accordance with the long-standing generalization that all metals conduct heat and electricity better than most non-metals. Oxidation states The inert s-pair effect is significant in the group-13 elements, especially the heavier ones like thallium. This results in a variety of oxidation states. In the lighter elements, the +3 state is the most stable, but the +1 state becomes more prevalent with increasing atomic number, and is the most stable for thallium. Boron is capable of forming compounds with lower oxidization states, of +1 or +2, and aluminium can do the same. Gallium can form compounds with the oxidation states +1, +2 and +3. Indium is like gallium, but its +1 compounds are more stable than those of the lighter elements. The strength of the inert-pair effect is maximal in thallium, which is only stable in the oxidation state of +1, although the +3 state is seen in some compounds. Periodic trends There are several trends that one could notice as they look at the properties of Boron group members. The Boiling Points of these elements drop from period to period, while densities tend to rise. 3

Element Boiling Point (C) Density (g/cm )

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Boron

4,000°

2.46

Aluminium 2,519°

2.7

Gallium

2,204°

5.904

Indium

2,072°

7.31

Thallium

1,473°

11.85

Nuclear With the exception of the synthetic ununtrium, all of the elements of the boron group have stable isotopes. Because all their atomic numbers are odd, boron, gallium and thallium have 10 only two stable isotopes, while aluminium and indium are monoisotopic, having only one. B 11 27 69 71 113 203 205 and B are both stable, as are Al, Ga and Ga, In, and Tl and Tl. All of these isotopes are readily found in macroscopic quantities in nature. In theory, though, all isotopes with an atomic number greater than 40 are supposed to be unstable to such decay modes as spontaneous fission and alpha decay. Conversely, all isotopes whose atomic numbers are less than 40 are theoretically supposed to be energetically stable to all forms of decay (with the exception of proton decay, which has never been observed). Like all other elements, the elements of the boron group have radioactive isotopes, either found in trace quantities in nature or produced synthetically. The longest-lived of these 115 14 unstable isotopes is the indium isotope In, with its extremely long half-life of 4.41 × 10 y. 7 This isotope is relatively important among indium's radioisotopes. The shortest-lived is B, −24 with a half-life of a mere 350±50 × 10 s, being the boron isotope with the fewest neutrons and a half-life long enough to measure. Some radioisotopes have important roles in scientific research; a few are used in the production of goods for commercial use or, more rarely, as a component of finished products. History The boron group has had many names over the years. According to former conventions it was Group IIIB in the European naming system and Group IIIA in the American. The group has also gained two collective names, "earth metals" and "triels". The latter name is derived from the Latin prefix tri- ("three") and refers to the three valence electrons that all of these elements, without exception, have in their valence shells. Boron was known to the ancient Egyptians, but only in the mineral borax. The metalloid element was not known in its pure form until 1808, when Humphry Davy was able to extract it by the method of electrolysis. Davy devised an experiment in which he dissolved a boroncontaining compound in water and sent an electric current through it, causing the elements of the compound to separate into their pure states. To produce larger quantities he shifted from electrolysis to reduction with sodium. Davy named the element boracium. At the same time two French chemists, Joseph Louis Gay-Lussac and Louis Jacques Thénard, used iron to reduce boric acid. The boron they produced was oxidized to boron oxide. Aluminium, like boron, was first known in minerals before it was finally extracted from alum, a common mineral in some areas of the world. Antoine Lavoisier and Humphry Davy had each separately tried to extract it. Although neither succeeded, Davy had given the metal its

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current name. It was only in 1825 that the Danish scientist Hans Christian Ørsted successfully prepared a rather impure form of the element. Many improvements followed, a significant advance being made just two years later by Friedrich Wöhler, whose slightly modified procedure still yielded an impure product. The first pure sample of aluminium is credited to Henri Etienne Sainte-Claire Deville, who substituted sodium for potassium in the procedure. At that time aluminium was considered precious, and it was displayed next to such metals as gold and silver. The method used today, electrolysis of aluminium oxide dissolved in cryolite, was developed by Charles Martin Hall and Paul Héroult in the late 1880s.

The mineral zinc blende, more commonly known as sphalerite, in which indium can occur. Thallium, the heaviest stable element in the boron group, was discovered by William Crookes and Claude-Auguste Lamy in 1861. Unlike gallium and indium, thallium had not been predicted by Dmitri Mendeleev, having been discovered before Mendeleev invented the periodic table. As a result, no one was really looking for it until the 1850s when Crookes and Lamy were examining residues from sulfuric acid production. In the spectra they saw a completely new line, a streak of deep green, which Crookes named after the Greek word θαλλόρ (thallos), referring to a green shoot or twig. Lamy was able to produce larger amounts of the new metal and determined most of its chemical and physical properties. Indium is the fourth element of the boron group but was discovered before the third, gallium, and after the fifth, thallium. In 1863 Ferdinand Reich and his assistant, Hieronymous Theodor Richter, were looking in a sample of the mineral zinc blende, also known as sphalerite (ZnS), for the spectroscopic lines of the newly discovered element thallium. Reich heated the ore in a coil of platinum metal and observed the lines that appeared in a spectroscope. Instead of the green thallium lines that he expected, he saw a new line of deep indigo-blue. Concluding that it must come from a new element, they named it after the characteristic indigo color it had produced. Gallium minerals were not known before August 1875, when the element itself was discovered. It was one of the elements that the inventor of the periodic table, Dmitri Mendeleev, had predicted to exist six years earlier. While examining the spectroscopic lines in zinc blende the French chemist Paul Emile Lecoq de Boisbaudran found indications of a new element in the ore. In just three months he was able to produce a sample, which he purified by dissolving it in a potassium hydroxide (KOH) solution and sending an electric current through it. The next month he presented his findings to the French Academy of Sciences, naming the new element after the Greek name for Gaul, modern France. It can be argued that the last confirmed element in the boron group, ununtrium, was not really "discovered", but "created" or synthesized. The element's synthesis is credited jointly to the Dubna Joint Institute for Nuclear Research team in Russia and the Lawrence Livermore National Laboratory in the United States, though it was the Dubna team who successfully conducted the experiment in August 2003. Element 113 (ununtrium) was

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discovered in the decay chain of element 115, or ununpentium, which produced a few precious atoms of ununtrium or "eka-thallium". The results were published in January of the following year. Since then around 13 atoms have been synthesized and various isotopes characterized. Etymology The name "boron" comes from the Arabic word for the mineral borax,(‫قروب‬, boraq) which was known before boron was ever extracted. The "-on" suffix is thought to have been taken from "carbon"; so the name can regarded as a portmanteau of "borax" and "carbon". Aluminium was named by Humphry Davy in the early 1800s. It is derived from the Greek word alumen, meaning bitter salt, or the Latin alum, the mineral. Gallium is derived from the Latin Gallia, referring to France, the place of its discovery. Indium comes from the Latin word indicum, meaning indigo dye, and refers to the element's prominent indigo spectroscopic line. Thallium, like indium, is named after the Greek word for the color of its spectroscopic line: thallos, meaning a green twig or shoot. "Ununtrium" is a temporary name assigned by the IUPAC (see IUPAC nomenclature), derived from the Latin names of the digits in the number 113. Occurrence and abundance

Boron Boron, with its atomic number of 5, is a very light element. Almost never found free in nature, it is very low in abundance, composing only 0.001% (10 ppm) of the Earth's crust. It is known to occur in over a hundred different minerals and ores, however: the main source is borax, but it is also found in colemanite, boracite, kernite, tusionite, berborite and fluoborite. Major world miners and extractors of boron include the United States, Turkey, Argentina, China, Bolivia and Peru. Turkey is by far the most prominent of these, accounting for around 70% of all boron extraction in the world. The United States is second, most of its yield coming from the state of California. Aluminium Aluminium, in contrast to boron, is the most abundant metal in the Earth's crust, and the third most abundant element. It composes about 8.2% (82,000 ppm) of the Earth, surpassed only by oxygen and silicon. It is like boron, however, in that it is uncommon in nature as a free element. This is due to aluminium’s tendency to attract oxygen atoms, forming several aluminium oxides. Aluminium is now known to occur in nearly as many minerals as boron, including garnets, turquoises and beryls, but the main source is the ore bauxite. The world's leading countries in the extraction of aluminium are Ghana, Surinam, Russia and Indonesia, followed by Australia, Guinea and Brazil. Gallium Gallium is a relatively rare element in the Earth's crust and is not found in as many minerals as its lighter homologues. Its abundance on the Earth is a mere 0.0018% (18 ppm). Its production is very low compared to other elements, but has increased greatly over the years as extraction methods have improved. Gallium can be found as a trace in a variety of ores, including bauxite and sphalerite, and in such minerals as diaspore and germanite. Trace amounts have been found in coal as well. The gallium content is greater in a few minerals,

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including gallite (CuGaS2), but these are too rare to be counted as major sources and make negligible contributions to the world's supply. Indium Indium is another rare element in the boron group. Even less abundant than gallium at only 0.000005% (0.05 ppm), it is the 61st most common element in the earth's crust. Very few indium-containing minerals are known, all of them scarce: an example is indite. Indium is found in several zinc ores, but only in minute quantities; likewise some copper and lead ores contain traces. As is the case for most other elements found in ores and minerals, the indium extraction process has become more efficient in recent years, ultimately leading to larger yields. Canada is the world's leader in indium reserves, but both the United States and China have comparable amounts. Thallium

A small bundle of fiberglass Thallium is neither rare nor common in the Earth's crust, but falls somewhere in the middle. Its abundance is estimated to be 0.00006% (0.6 ppm). Thallium is the 56th most common element in the earth's crust, more abundant than indium by a sizeable amount. It is found on the ground in some rocks, in the soil and in clay. Many sulfide ores of iron, zinc and cobalt contain thallium. In minerals it is found in moderate quantities: some examples are crookesite (in which it was first discovered), lorandite, routhierite, bukovite, hutchinsonite and sabatierite. There are other minerals that contain small amounts of thallium, but they are very rare and do not serve as primary sources. Macedonia is a notable thallium extractor and producer.

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Ununtrium Ununtrium is an element that is never found in nature but has been created in a laboratory. It is therefore classified as a synthetic element with no stable isotopes. Applications With the exception of synthetic ununtrium, all the elements in the boron group have numerous uses and applications in the production and content of many items. Boron has found many industrial applications in recent decades, and new ones are still being found. A common application is in fiberglass. There has been rapid expansion in the market for borosilicate glass; most notable among its special qualities is a much greater resistance to thermal expansion than regular glass. Another commercially expanding use of boron and its derivatives is in ceramics. Several boron compounds, especially the oxides, have unique and valuable properties that have led to their substitution for other materials that are less useful. Boron may be found in pots, vases, plates, and ceramic pan-handles for its insulating properties. The compound borax is used in bleaches, for both clothes and teeth. The hardness of boron and some of its compounds give it a wide array of additional uses. A small part (5%) of the boron produced finds use in agriculture. Aluminium is a metal with numerous familiar uses in everyday life. It is most often encountered in construction materials, in electrical devices, especially as the conductor in cables, and in tools and vessels for cooking and preserving food. Aluminium's lack of reactivity with food products makes it particularly useful for canning. Its high affinity for oxygen makes it a powerful reducing agent. Finely powdered pure aluminium oxidizes rapidly in air, generating a huge amount of heat in the process (burning at about 5500 째F or 3037 째C), leading to applications in welding and elsewhere that a large amount of heat is needed. Aluminium is a component of alloys used for making lightweight bodies for aircraft. Cars also sometimes incorporate aluminium in their framework and body, and there are similar applications in military equipment. Less common uses include components of decorations and some guitars. The element is also sees use in a diverse range of electronics.

Gallium is one of the chief components of blue LEDs Gallium and its derivatives have only found applications in recent decades. Gallium arsenide has been used in semiconductors, in amplifiers, in solar cells (for example in satellites) and in tunnel diodes for FM transmitter circuits. Gallium alloys are used mostly for dental purposes. Gallium ammonium chloride is used for the leads in transistors. A major application of gallium is in LED lighting. The pure element has been used as a dopant in [citation needed] semiconductors, and has additional uses in electronic devices with other elements. Gallium has the property of being able to 'wet' glass and porcelain, and thus can

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be used to make mirrors and other highly reflective objects. Gallium can be added to alloys of other metals to lower their melting points. Indium's uses can be divided into four categories: the largest part (70%) of the production is used for coatings, usually combined as indium tin oxide (ITO); a smaller portion (12%) goes into alloys and solders; a similar amount is used in electrical components and in semiconductors; and the final 6% goes to minor applications. Among the items in which indium may be found are platings, bearings, display devices, heat reflectors, phosphors, and nuclear control rods. Indium tin oxide has found a wide range of applications, including glass coatings, solar panels, streetlights, electrophosetic displays (EPDs), electroluminescent displays (ELDs), plasma display panels (PDPs), electrochemic displays (ECs), field emission displays (FEDs), sodium lamps, windshield glass and cathode ray tubes, making it the single most important indium compound. Thallium is used in its elemental form more often than the other boron-group elements. Uncompounded thallium is used in low-melting glasses, photoelectric cells, switches, mercury alloys for low-range glass thermometers, and thallium salts. It can be found in lamps and electronics, and is also used in myocardial imaging. The possibility of using thallium in semiconductors has been researched, and it is a known catalyst in organic synthesis. Thallium hydroxide (TlOH) is used mainly in the production of other thallium compounds. Thallium sulfate (Tl2SO4) is an outstanding vermin-killer, and it is a principal component in some rat and mouse poisons. However, the United States and some European countries have banned the substance because of its high toxicity to humans. In other countries, though, the market for the substance is growing. Tl2SO4 is also used in optical systems. Biological role None of the group-13 elements has a major biological role in complex animals, but some are at least associated with a living being. As in other groups, the lighter elements usually have more biological roles than the heavier. The heaviest ones are toxic, as are the other elements in the same periods. Boron is essential in most plants, whose cells use it for such purposes as strengthening cell walls. It is found in humans, certainly as a trace element, but there is ongoing debate over its significance in human nutrition. Boron's chemistry does allow it to form complexes with such important molecules as carbohydrates, so it is plausible that it could be of greater use in the human body than previously thought. Boron has also been shown to be able to replace iron in some of its functions, particularly in the healing of wounds. Aluminium has no known biological role in plants or animals. Gallium is not essential for the human body, but its relation to iron(III) allows it to become bound to proteins that transport and store iron. Gallium can also stimulate metabolism. Indium and its heavier homologues have no biological role, although indium salts in small doses, like gallium, can stimulate metabolism. Toxicity All of the elements in the boron group can be toxic, given a high enough dose. Some of them are only toxic to plants, some only to animals, and some to both. As an example of boron toxicity, it has been observed to harm barley in concentrations exceeding 20 mM. The symptoms of boron toxicity are numerous in plants, complicating research: they include reduced cell division, decreased shoot and root growth, decreased production of leaf chlorophyll, inhibition of photosynthesis, lowering of stomata conductance, reduced proton extrusion from roots, and deposition of lignin and suborgin. Aluminium does not present a prominent toxicity hazard in small quantities, but very large doses are slightly toxic. Gallium is not considered toxic, although it may have some minor

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effects. Indium is not toxic and can be handled with nearly the same precautions as gallium, but some of its compounds are slightly to moderately toxic. Thallium, unlike gallium and indium, is extremely toxic, and has caused many poisoning deaths. Its most noticeable effect, apparent even from tiny doses, is hair loss all over the body, but it causes a wide range of other symptoms, disrupting and eventually halting the functions of many organs. The nearly colorless, odorless and tasteless nature of thallium compounds has led to their use by murderers. The incidence of thallium poisoning, intentional and accidental, increased when thallium (with its similarly toxic compound, thallium sulfate) was introduced to control rats and other pests. The use of thallium pesticides has therefore been prohibited since 1975 in many countries, including the USA. Ununtrium is a highly unstable element and decays by emitting alpha particles. Due to its strong radioactivity, it would definitely be extremely toxic, although significant quantities of ununtrium (larger than a few atoms) have not yet been assembled. Bibliography ď&#x201A;ˇ

Downs, Anthony John (1993). Chemistry of aluminium, gallium, indium, and thallium. Chapman and Hall Inc. pp. 197â&#x20AC;&#x201C;201. ISBN 978-0-7514-0103-5.

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Chapter 3: Carbon group

Carbon group (group 14)

boron group ←

→ pnictogens

IUPAC group number

Name by element

carbon group

Trivial name

tetrels, crystallogens

CAS group number (US)

IVA

old IUPAC number (European)IVB

↓ Period 2

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Carbon

(C)

6 Polyatomic nonmetal 3

Silicon

(Si)

14 Metalloid 4

Germanium

(Ge)

32 Metalloid 5

Tin

(Sn)

50 Poor metal

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Lead

(Pb)

82 Poor metal 7 Flerovium

(Fl)

114 unknown chemical properties

Legend primordial element synthetic element Atomic number color: black=solid

The carbon group is a periodic table group consisting of carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb), and flerovium (Fl). In modern IUPAC notation, it is called Group 14. In the field of semiconductor physics, it is still universally called Group IV. The group was once also known as the tetrels (from Greek tetra, four), stemming from the Roman numeral IV in the group names, or (not coincidentally) from the fact that these elements have four valence electrons (see below). The group is sometimes also referred to as tetragens or crystallogens. Characteristics

Chemical Like other groups, the members of this family show patterns in electron configuration, especially in the outermost shells, resulting in trends in chemical behavior:

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Element

No. of electrons/shell

Carbon

2, 4

14 Silicon

2, 8, 4

32 Germanium 2, 8, 18, 4 50 Tin

2, 8, 18, 18, 4

82 Lead

2, 8, 18, 32, 18, 4

114 Flerovium

2, 8, 18, 32, 32, 18, 4 (predicted)

Each of the elements in this group has 4 electrons in its outer energy level. The last orbital of 2 all these elements is the p orbital. In most cases, the elements share their electrons. The tendency to lose electrons increases as the size of the atom increases, as it does with 4â&#x2C6;&#x2019; increasing atomic number. Carbon alone forms negative ions, in the form of carbide (C ) ions. Silicon and germanium, both metalloids, each can form +4 ions. Tin and lead both are metals while flerovium is a synthetic, radioactive (its half life is very short), element that may have a few noble gas-like properties, though it is still most likely a post-transition metal. Tin and lead are both capable of forming +2 ions. Carbon forms tetrahalides with all the halogens except astatine. Carbon also forms three oxides: carbon monoxide, carbon suboxide (C3O2), and carbon dioxide. Carbon forms disulfides and diselenides. Silicon forms two hydrides: SiH4 and Si2H6. Silicon forms tetrahalides with fluorine, chlorine, and iodine. Silicon also forms a dioxide and a disulfide.Silicon nitride has the formula Si3N4. Germanium forms two hydrides: GeH4 and Ge2H6. Germanium forms tetrahalides with all halogens except astatine and forms dihalides with all halogens except bromine and astatine. Germanium bonds to all natural single chalcogens except polonium, and forms dioxides, disulfides, and diselenides. Germanium nitride has the formula Ge 3N4. Tin forms two hydrides: SnH4 and Sn2H6. Tin forms dihalides and tetrahalides with all halogens except astatine. Tin forms chalcogenides with one of each naturally occurring chalcogen except polonium, and forms chalcogenides with two of each naturally occurring chalcogen except polonium and tellurium. Lead forms one hydride, which has the formula PbH4. Lead forms dihalides and tetrahalides with fluorine and chlorine, and forms a tetrabromide and a lead diiodide. Lead forms four oxides, a sulfide, a selenide, and a telluride. There are no known compounds of flerovium.

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Physical The boiling points of the carbon group tend to get lower with the heavier elements. Carbon, the lightest carbon group element, sublimates at 3825 °C. Silicon's boiling point is 3265 °C, germanium's is 2833 °C, tin's is 2602 °C, and lead's is 1749 °C. The melting points of the carbon group elements have roughly the same trend as their boiling points. Silicon melts at 1414 °C, germanium melts at 939 °C, tin melts at 232 °C, and lead melts at 328 °C. Carbon's crystal structure is hexagonal. Silicon and germanium have face-centered cubic crystal structures. Tin has a tetragonal crystal structure. Lead has a face-centered cubic crystal structure. The densities of the carbon group elements tend to increase with increasing atomic number. Carbon has a density of 2.26 grams per cubic centimeter, silicon has a density of 2.33 grams per cubic centimeter, germanium has a density of 5.32 grams per cubic centimeter. Tin has a density of 7.26 grams per cubic centimeter, and lead has a density of 11.3 grams per cubic centimeter. The atomic radii of the carbon group elements tend to increase with increasing atomic number. Carbon's atomic radius is 77 picometers, silicon's is 118 picometers, germanium's is 123 picometers, tin's is 141 picometers, and lead's is 175 picometers. Allotropes Main article: Allotropes of carbon Carbon has multiple allotropes. The most common is graphite, which is carbon in the form of stacked sheets. Another form of carbon is diamond, but this is relatively rare. Amorphous carbon is a third allotrope of carbon; it is a component of soot. Another allotrope of carbon is a fullerene, which has the form of sheets of carbon atoms folded into a sphere. A fifth allotrope of carbon, discovered in 2003, is called graphene, and is in the form of a layer of carbon atoms arranged in a honeycomb-shaped formation. Silicon has two known allotropes that exist at room temperature. These allotropes are known as the amorphous and the crystalline allotropes. The amorphous allotrope is a brown powder. The crystalline allotrope is gray and has a metallic luster. Tin has two allotropes: α-tin, also known as gray tin, and β-tin. Tin is typically found in the βtin form, a silvery metal. However, at standard pressure, β-tin converts to α-tin, a gray powder, at temperatures below 56° Fahrenheit. This can cause tin objects in cold temperatures to crumble to gray powder in a process known as tin rot. Nuclear At least two of the carbon group elements (tin and lead) have magic nuclei, meaning that these elements are more common and more stable than elements that do not have a magic nucleus. Isotopes There are 15 known isotopes of carbon. Of these, three are naturally occurring. The most common is stable carbon-12, followed by stable carbon-13.Carbon-14 is a natural radioactive isotope with a half-life of 5,730 years.

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23 isotopes of silicon have been discovered. Five of these are naturally occurring. The most common is stable silicon-28, followed by stable silicon-29 and stable silicon-30. Silicon-32 is a radioactive isotope that occurs naturally as a result of radioactive decay of actinides, and via spallation in the upper atmosphere. Silicon-34 also occurs naturally as the result of radioactive decay of actinides. 32 isotopes of germanium have been discovered. Five of these are naturally occurring. The most common is the stable isotope germanium-74, followed by the stable isotope germanium-72, the stable isotope germanium-70, and the stable isotope germanium-73. The isotope germanium-76 is a primordial radioisotope. 40 isotopes of tin have been discovered. 14 of these occur in nature. The most common is the stable isotope tin-120, followed by the stable isotope tin-118, the stable isotope tin-116, the stable isotope tin-119, the stable isotope tin-117, the primordial radioisotope tin-124, the stable isotope tin-122, the stable isotope tin-112, and the stable isotope tin-114. Tin also has four radioisotopes that occur as the result of the radioactive decay of uranium. These isotopes are tin-121, tin-123, tin-125, and tin-126. 38 isotopes of lead have been discovered. 9 of these are naturally occurring. The most common isotope is the primordial radioisotope lead-208, followed by the primordial radioisotope lead-206, the primordial radioisotope lead-207, and the primordial radioisotope lead-204. 4 isotopes of lead occur from the radioactive decay of uranium and thorium. These isotopes are lead-209, lead-210, lead-211, and lead-212. 4 isotopes of flerovium (flerovium-286, flerovium-287, flerovium-288, and flerovium-289) have been discovered. None of these are naturally occurring. Flerovium's most stable isotope is flerovium-289, which has a half-life of 2.6 seconds. Occurrence Carbon accumulates as the result of stellar fusion in most stars, even small ones. Carbon is present in the earth's crust in concentrations of 480 parts per million, and is present in seawater at concentrations of 28 parts per million. Carbon is present in the atmosphere in the form of carbon monoxide, carbon dioxide, and methane. Carbon is a key constituent of carbonate minerals, and is in hydrogen carbonate, which is common in seawater. Carbon forms 22.8% of a typical human. Silicon is present in the earth's crust at concentrations of 28%, making it the second most abundant element there. Silicon's concentration in seawater can vary from 30 parts per billion on the surface of the ocean to 2000 parts per billion deeper down. Silicon dust occurs in trace amounts in earth's atmosphere. Silicate minerals are the most common type of mineral on earth. Silicon makes up 14.3 parts per million of the human body on average. Only the largest stars produce silicon via stellar fusion. Germanium makes up 2 parts per million of the earth's crust, making it the 52nd most abundant element there. On average, germanium makes up 1 part per million of soil. Germanium makes up 0.5 parts per trillion of seawater. Organogermanium compounds are also found in seawater. Germanium occurs in the human body at concentrations of 71.4 parts per billion. Germanium has been found to exist in some very faraway stars. Tin makes up 2 parts per million of the earth's crust, making it the 49th most abundant element there. On average, tin makes up 1 part per million of soil. Tin exists in seawater at concentrations of 4 parts per trillion. Tin makes up 428 parts per million of the human body. Tin (IV) oxide occurs at concentrations of 0.1 to 300 parts per million in soils. Tin also occurs in concentrations of one part per thousand in igneous rocks.

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Lead makes up 14 parts per million of the earth's crust, making it the 36th most abundant element there. On average, lead makes up 23 parts per million of soil, but the concentration can reach 20000 parts per million (2 percent) near old lead mines. Lead exists in seawater at concentrations of 2 parts per trillion. Lead makes up 0.17% of the human body by weight. Human activity releases more lead into the environment than any other metal. Flerovium only occurs in particle accelerators. History

Antiquity discoveries Carbon, tin, and lead are a few of the elements well known in the ancient world—together with sulfur, iron, copper, mercury, silver, and gold. Carbon as an element was discovered by the first human to handle charcoal from a fire. Modern carbon chemistry dates from the development of coals, petroleum, and natural gas as fuels and from the elucidation of synthetic organic chemistry, both substantially developed since the 19th century. The origins of tin seem to be lost in history. It appears that bronzes, which are alloys of copper and tin, were used by prehistoric man some time before the pure metal was isolated. Bronzes were common in early Mesopotamia, the Indus Valley, Egypt, Crete, Israel, and Peru. Much of the tin used by the early Mediterranean peoples apparently came from the Scilly Isles and Cornwall in the British Isles, where mining of the metal dates from about 300–200 BCE. Tin mines were operating in both the Inca and Aztec areas of South and Central America before the Spanish conquest. Lead is mentioned often in early Biblical accounts. The Babylonians used the metal as plates on which to record inscriptions. The Romans used it for tablets, water pipes, coins, and even cooking utensils; indeed, as a result of the last use, lead poisoning was recognized in the time of Augustus Caesar. The compound known as white lead was apparently prepared as a decorative pigment at least as early as 200 BCE. Modern developments date to the exploitation in the late 18th century of deposits in the Missouri–Kansas–Oklahoma area in the United States. Modern discoveries Amorphous elemental silicon was first obtained pure in 1824 by the Swedish chemist Jöns Jacob Berzelius; impure silicon had already been obtained in 1811. Crystalline elemental silicon was not prepared until 1854, when it was obtained as a product of electrolysis. In the form of rock crystal, however, silicon was familiar to the predynastic Egyptians, who used it for beads and small vases; to the early Chinese; and probably to many others of the ancients. The manufacture of glass containing silica was carried out both by the Egyptians — at least as early as 1500 BCE — and by the Phoenicians. Certainly, many of the naturally occurring compounds called silicates were used in various kinds of mortar for construction of dwellings by the earliest people. Germanium is one of three elements the existence of which was predicted in 1869 by the Russian chemist Dmitri Mendeleev when he first devised his periodic table. However, the element was not actually discovered for some time. In September 1885, a miner discovered a mineral sample in a silver mine and gave it to the mine manager, who determined that it was a new mineral and sent the mineral to Clemens A. Winkler. Winkler realized that the

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sample was 75% silver, 18% sulfur, and 7% of an undiscovered element. After several months, Winkler isolated the element and determined that it was element 32. The first attempt to discover flerovium (then referred to as "element 114") was in 1969, at the Joint Institute for Nuclear Research, but it was unsuccessful. In 1977, researchers at the Joint Institute for Nuclear Research bombarded plutonium-244 atoms with calcium-48, but were again unsuccessful. This nuclear was repeated in 1998, this time successfully. Etymologies The word "carbon" comes from the Latin word carbo, meaning "charcoal".The word "silicon" comes from the Latin word silex or silicis, which mean "flint". The word "germanium" comes from the word germania, which is Latin for Germany, the county where germanium was discovered. The word "tin" derives from the Old English word tin. The word "lead" comes from the Old English word lead. Applications Carbon is most commonly used in its amorphous form. In this form, carbon is used for steelmaking, as carbon black, as a filling in tires, in respirators, and as activated charcoal. Carbon is also used in the form of graphite is commonly used as the lead in pencils. Diamond, another form of carbon, is commonly used in jewelery.Carbon fibers are used in numerous applications, such as satellite struts, because the fibers are highly strong yet elastic. Silicon dioxide has a wide variety of applications, including toothpaste, construction fillers, and silica is a major component of glass. 50% of pure silicon is devoted to the manufacture of metal alloys. 45% of silicon is devoted to the manufacture of silicones. Silicon is also commonly used in semiconductors since the 1950s. Germanium was used in semiconductors until the 1950s, when it was replaced by silicon. Radiation detectors contain germanium. Germanium oxide is used in fiber optics and wideangle camera lenses. A small amount of germanium mixed with silver can make silver tarnish-proof. The resulting alloy is known as argentium. Solder is the most important use of tin; 50% of all tin produced goes into this application. 20% of all tin produced is used in tin plate. 20% of tin is also used by the chemical industry. Tin is also a constituent of numerous alloys, including pewter. Tin (IV) oxide has been commonly used in ceramics for thousands of years. Cobalt stannate is a tin compound which is used as a cerulean blue pigment. 80% of all lead produced goes into lead-acid batteries. Other applications for lead include weights, pigments, and shielding against radioactive materials. Lead was historically used in gasoline in the form of tetraethyl lead, but this application has been discontinued due to concerns of toxicity. Production Carbon's allotrope diamond is produced mostly by Russia, Botswana, Congo, Canada, and South Africa. 80% of all synthetic diamonds are produced by Russia. China produces 70% of the world's graphite. Other graphite-mining countries are Brazil, Canada, and Mexico. Silicon can be produced by heating silica with carbon.

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There are some germanium ores, such as germanite, but these are not mined on account of being rare. Instead, germanium is extracted from the ores of metals such as zinc. In Russia and China, germanium is also separated from coal deposits. Germanium-containing ores are first treated with chlorine to form germanium tetrachloride, which is mixed with hydrogen gas. Then the germanium is further refined by zone refining Roughly 140 metric tons of germanium are produced each year. Mines output 300,000 metric tons of tin each year. China, Indonesia, Peru, Bolivia, and Brazil are the main producers of tin. The method by which tin is produced is to head the tin mineral cassiterite (SnO2) with coke. The most commonly mined lead ore is galena (lead sulfide). 4 million metric tons of lead are newly mined each year, mostly in China, Australia, the United States, and Peru. The ores are mixed with coke and limestone and roasted to produce pure lead. Most lead is recycled from lead batteries. The total amount of lead ever mined by humans amounts to 350 million metric tons. Biological role Carbon is a key element to all known life. It is in all organic compounds, for example, DNA, steroids, and proteins. Carbon's importance to life is primarily due to its ability to form numerous bonds with other elements. There are 16 kilograms of carbon in a typical 70kilogram human. Silicon-based life's feasibility is commonly discussed. However, it is less able than carbon to form elaborate rings and chains. Silicon in the form of silicon dioxide is used by diatoms and sea sponges to form their cell walls and skeletons. Silicon is essential for bone growth in chickens and rats and may also be essential in humans. Humans consume on average between 20 and 1200 milligrams of silicon per day, mostly from cereals. There is 1 gram of silicon in a typical 70-kilogram human. A biological role for germanium is not known, although it does stimulate metabolism. In 1980, germanium was reported by Kazuhiko Asai to benefit health, but the claim has not been proven. Some plants take up germanium from the soil in the form of germanium oxide. These plants, which include grains and vegetables contain roughly 0.05 parts per million of germanium. The estimated human intake of germanium is 1 milligram per day. There are 5 milligrams of germanium in a typical 70-kilogram human. Tin has been shown to be essential for proper growth in rats, but there is, as of 2013, no evidence to indicate that humans need tin in their diet. Plants do not require tin. However, plants do collect tin in their roots. Wheat and corn contain seven and three parts per million respectively. However, the level of tin in plants can reach 2000 parts per million if the plants are near a tin smelter. On average, humans consume 0.3 milligrams of tin per day. There are 30 milligrams of tin in a typical 70-kilogram human. Lead has no known biological role, and is in fact highly toxic, but some microbes are able to survive in lead-contaminated environments. Some plants, such as cucumbers contain up to tens of parts per million of lead. There are 120 milligrams of lead in a typical 70-kilogram human. Toxicity Elemental carbon is not generally toxic, but many of its compounds are, such as carbon monoxide and hydrogen cyanide. However, carbon dust can be dangerous because it lodges in the lungs in a manner similar to asbestos.

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Silicon minerals are not typically poisonous. However, silicon dioxide dust, such as that emitted by volcanoes can cause adverse health effects if it enters the lungs. Germanium can interfere with such enzymes as lactate and alcohol dehydrogenase. Organic germanium compounds are less toxic than inorganic germanium compounds. Germanium has a low degree of oral toxicity in animals. Severe germanium poisoning can cause death by respiratory paralysis. Some tin compounds are toxic to ingest, but most inorganic compounds of tin are considered nontoxic. Organic tin compounds, such as trimethyl tin and triethyl tin are highly toxic, and can disrupt metabolic processes inside cells. Lead and its compounds, such as lead acetate are highly toxic. Lead poisoning can cause headaches, stomach pain, constipation, and gout.

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Chapter 4: Pnictogen Pnictogens

carbon group ←

→ chalcogens

IUPAC group number

Name by element

nitrogen group

Trivial name

pnictogens

CAS group number (US)

old IUPAC number (European)VB

↓ Period

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Nitrogen

(N)

7 Diatomic nonmetal 3

Phosphorus

(P)

15 Polyatomic nonmetal 4

Arsenic

(As)

33 Metalloid

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Antimony

(Sb)

51 Metalloid 6

Bismuth

(Bi)

83 Poor metal 7 Ununpentium

(Uup)

115 unknown chemical properties

Legend primordial element synthetic element Atomic number color: red=gas, black=solid

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The pnictogens/ˈnɪktədʒɨnz/ are the chemical elements in group 15 of the periodic table. This group is also known as the nitrogen family. It consists of the elements nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), bismuth (Bi) and the synthetic element ununpentium (Uup) (unconfirmed). In modern IUPAC notation, it is called Group 15. In CAS and the old IUPAC systems it was called Group VA and Group VB, respectively (pronounced "group five A" and "group five B", "V" for the Roman numeral 5). In the field of semiconductor physics, it is still usually called Group V. The "five" ("V") in the historical names comes from the "pentavalency" of nitrogen, reflected by the stoichiometry of compounds such as N 2O5. Characteristics

Chemical Like other groups, the members of this family show patterns in electron configuration, especially in the outermost shells, resulting in trends in chemical behavior: Z Element

No. of electrons/shell

2, 5

nitrogen

15 phosphorus 2, 8, 5 33 arsenic

2, 8, 18, 5

51 antimony

2, 8, 18, 18, 5

83 bismuth

2, 8, 18, 32, 18, 5

This group has the defining characteristic that all the component elements have 5 electrons in their outermost shell, that is 2 electrons in the s subshell and 3 unpaired electrons in the p subshell. They are therefore 3 electrons short of filling their outermost electron shell in their non-ionized state. The most important elements of this group are nitrogen (N), which in its diatomic form is the principal component of air, and phosphorus (P), which, like nitrogen, is essential to all known forms of life. Compounds Binary compounds of the group can be referred to collectively as pnictides. The spelling derives from the Greek verb πνίγειν (pnígein), to "choke" or "stifle", which is a property of molecular nitrogen in the absence of oxygen; it can also be used as a mnemonic for the two most common members, P and N. The name pentels (from Greek πέντε, pénte, five) was also used for this group at one time, stemming from the earlier group naming convention (Group VB). Pnictide compounds tend to be exotic. Various properties that some pnictides have include being dimagnetic and paramagnetic at room temperature, being transparent, and generating electricity when heated. Other pnictides include the ternary rare-earth main-group variety of

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pnictides. These are in the form of REaMbPnc, where M is a carbon group or boron group element and Pn is any pnictogen except nitrogen. These compounds are between ionic and covalent compounds and thus have unusual bonding properties. These elements are also noted for their stability in compounds due to their tendency for forming double and triple covalent bonds. This is the property of these elements which leads to their potential toxicity, most evident in phosphorus, arsenic and antimony. When these substances react with various chemicals of the body, they create strong free radicals not easily processed by the liver, where they accumulate. Paradoxically it is this strong bonding which causes nitrogen and bismuth's reduced toxicity (when in molecules), as these form strong bonds with other atoms which are difficult to split, creating very unreactive molecules. For example N2, the diatomic form of nitrogen, is used as an inert gas in situations where using argon or another noble gas would be too expensive. The upper pnictogens, that is, nitrogen, phosphorus, and arsenic tend to form −3 charges. Antimony and bismuth can either take on a +3 or +5, by losing its p-shell electrons or losing its p-shell and s-shell electrons, respectively. Physical The pnictogens consist of two nonmetals (one gas, one solid), two metalloids, one metal, and one element with unknown chemical properties. All the elements in the group are solids at room temperature, except for nitrogen which is gaseous at room temperature. Nitrogen and bismuth, despite both being pnictogens, are very different in their physical properties. For instance, at STP nitrogen is a transparent nonmetallic gas, while bismuth is a silverywhite metal. The densities of the pnictogens increase towards the heavier pnictogens. Nitrogen's density is 0.001251 grams per cubic centimeter at STP. Phosphorus's density is 1.82 grams per cubic centimeter at STP, arsenic's is 5.72 grams per cubic centimeter, antimony's is 6.68 grams per cubic centimeter, and bismuth's is 9.79 grams per cubic centimeter. Nitrogen's melting point is -210°C and its boiling point is -196°C. Phosphorus's melting point is 44°C and its boiling point is 280°C. Arsenic is one of only two elements to sublimate at standard pressure; it does this at 603°C. Antimony's melting point is 631°C and its boiling point is 1587°C. Bismuth's melting point is 271°C and its boiling point is 1564°C. Nitrogen's crystal structure is hexagonal. Phosphorus's crystal structure is cubic. Arsenic, antimony, and bismuth all have rhombohedral crystal structures. History The nitrogen compound sal ammoniac (ammonium chloride) was known since the time of the Ancient Egyptians. In the 1760s two scientists, Henry Cavendish and Joseph Priestley, isolated nitrogen from air, but neither realized the presence of an undiscovered element. It was not until several years later, in 1772, that Daniel Rutherford realized that the gas was indeed nitrogen. The scientist Hennig Brandt first discovered phosphorus in Hamburg in 1669. Brandt produced the element by heating evaporated urine and condensing the resulting phosphorus vapor in water. Brandt initially thought that he had discovered the Philosopher's Stone, but eventually realized that this was not the case.

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Arsenic compounds have been known for at least 5000 years, and the ancient Greek Theophrastus recognized the arsenic minerals called realgar and orpiment. Elemental arsenic was discovered in the 13th century by Albertus Magnus. Antimony was well-known to the ancients. A 5000-year-old vase made of nearly pure antimony exists in the Louvre. Antimony compounds were used in dyes in the Babylonian times. The antimony mineral stibnite may have been a component of Greek fire. Bismuth was first discovered by an alchemist in 1400. Within 80 years of bismuth's discovery, it had applications in printing and decorated caskets. The Incas were also using bismuth in knives by 1500. Bismuth was originally thought to be the same as lead, but in 1753, Claude-François Geoffrey proved that bismuth was different from lead. In 1977, the Joint Institute for Nuclear Research attempted to produce ununpentium by bombarding plutonium-244 atoms with calcium-40 atoms, but were unsuccessful. Ununpentium was successfully produced in 2003 by bombarding americium-243 atoms with calcium-48 atoms. Etymology The term "pnictogen" was suggested by the Dutch chemist Anton Eduard van Arkel in the early 1950s. It is also spelled "pnicogen" or "pnigogen". The term "pnicogen" is rarer than the term "pnictogen", and the ratio of research papers using "pnictogen" to those using "pnicogen" is 2.5 to 1. It comes from the Greek root πνιγ- (choke, strangle), and thus the word "pnictogen" is also a reference to the Dutch and German names for nitrogen (stikstof, Stickstoff, "suffocating substance", i.e. portion of air unsuitable for breathing). Hence, "pnictogen" could be translated as "suffocator maker". The word "pnictide" also comes from the same root. Occurrence

A collection of nitrogen-group chemical element samples. Nitrogen makes up 25 parts per million of the earth's crust, 5 parts per million of soil on average, 100 to 500 parts per trillion of seawater, and 78% of dry air. The majority of nitrogen on earth is in the form of nitrogen gas, but some nitrate minerals do exist. Nitrogen makes up 2.5% of a typical human by weight. Phosphorus makes up 0.1% of the earth's crust, making it the 11th most abundant element there. Phosphorus makes up 0.65 parts per million of soil, and 15 to 60 parts per billion of seawater. There are 200 million metric tons of accessible phosphates on earth. Phosphorus makes up 1.1% of a typical human by weight.

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Arsenic makes up 1.5 parts per million of the earth's crust, making it the 53rd most abundant element there. The soils contain 1 to 10 parts per million of arsenic, and seawater contains 1.6 parts per billion of arsenic. Arsenic makes up 100 parts per billion of a typical human by weight. Some arsenic exists in elemental form, but most arsenic is found in the arsenic minerals orpiment, realgar, arsenopyrite, and enargite. Antimony makes up 0.2 parts per million of the earth's crust, making it the 63rd most abundant element there. The soils contain 1 part per million of antimony on average, and seawater contains 300 parts per trillion of antimony on average. A typical human contains 28 parts per billion of antimony by weight. Some elemental antimony occurs in silver deposits. Bismuth makes up 48 parts per billion of the earth's crust, making it the 70th most abundant element there. The soils contain approximately 0.25 parts per million of bismuth, and seawater contains 400 parts per trillion of bismuth. Bismuth most commonly occurs as the mineral bismuthinite, but bismuth also occurs in elemental form or in sulfide ores. Ununpentium is produced several atoms at a time in particle accelerators. Production

Nitrogen Nitrogen can be produced by fractional distillation of air. Nitrogen can also be produced in a large scale by burning hydrocarbons or hydrogen in air. On a smaller scale, it is also possible to make nitrogen by heating barium azide. Additionally, the following reactions produce nitrogen: NH4 + NO2 = N2 + 2H2O +

8NH3 + 3Br2 = N2 + 6NH4 + 6Br

â&#x2C6;&#x2019;

2NH3 + 3CuO = N2 + 3H2O + 2Cu

Phosphorus There are two principal methods for producing phosphorus. One is to mix crushed phosphate rocks with phosphoric or sulfuric acid, producing calcium hydrogen phosphates. The other is to reduce phosphates with carbon in an electric furnace. Arsenic Arsenic is mostly produced in Sweden. Most arsenic is prepared by heating the mineral arsenopyrite in the presence of air. This forms As 4O6, from which arsenic can be extracted via carbon reduction. However, it is also possible to make metallic arsenic by heating arsenopyrite at 650Â° to 700Â° Celsius without oxygen. Antimony With sulfide ores, the method by which antimony is produced depends on the amount of antimony in the raw ore. If the ore contains 25% to 45% antimony by weight, then crude

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antimony is produced by smelting the ore in a blast furnace. If the ore contains 45% to 60% antimony by weight, antimony is obtained by heating the ore, also known as liquidation. Ores with more than 60% antimony by weight are chemically displaced with iron shavings from the molten ore,resulting in impure metal. If an oxide ore of antimony contains less than 30% antimony by weight, the ore is reduced in a blast furnace. If the ore contains closer to 50% antimony by weight, the ore is instead reduced in a reverberatory furnace. Antimony ores with mixed sulfides and oxides are smelted in a blast furnace. Bismuth Bismuth minerals do occur, but it is more economic to produce bismuth as a by-product of lead. In China, bismuth is also found in tungsten and zinc ores. Ununpentium Ununpentium is produced a few atoms at a time in particle accelerators. Applications 

Liquid nitrogen is a commonly used cryogenic liquid.



Nitrogen in the form of ammonia a nutrient critical to most plants' survival.



Phosphorus is used in matches and incendiary bombs.



Phosphate fertilizer helps feed much of the world.



Arsenic was historically used as a Paris green pigment, but is not used this way anymore, due to arsenic's extreme toxicity.



Arsenic in the form of organoarsenic compounds is sometimes used in chicken feed



Antimony is alloyed with lead to produce some bullets.



Antimony currency was briefly used in the 1930s in parts of China, but this use was discontinued as antimony is soft and toxic



Bismuth subsalicylate is the active ingredient in Pepto-Bismol.

Biological role Nitrogen is a component of molecules critical to life on earth, such as DNA and amino acids. Nitrates occur in some plants, such as spinach and lettuce. A typical 70-kilogram human contains 1.8 kilograms of nitrogen. Phosphorus in the form of phosphates occur in compounds important to life, such as DNA and ATP. Humans typically consume 1 to 2 milligrams of phosphorus per day. Phosphorus is found in several kinds of fish, liver, turkey, chicken, and eggs. Phosphate deficiency is a

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problem known as hypophosphatemia. A typical 70-kilogram human contains 480 grams of phosphorus. Arsenic promotes growth in chickens and rats, and may be essential for humans. Arsenic has been shown to be helpful in metabolizing the amino acid arginine. There are 7 milligrams of arsenic in a typical 70-kilogram human. Antimony is not known to have a biological role. Plants take up only trace amounts of antimony. There are approximately 2 milligrams of antimony in a typical 70-kilogram human. Bismuth is not known to have a biological role. Humans ingest on average less than 20 micrograms of bismuth per day. There is less than 500 micrograms of bismuth in a typical 70-kilogram human. Toxicity Breathing in pure nitrogen gas is deadly, causing nitrogen asphyxiation, although nitrogen gas is completely nontoxic if there is also enough oxygen to breathe. The buildup of nitrogen bubbles in the blood, such as during deep-sea diving, can cause a condition known as the "Bends". Many nitrogen compounds, such as hydrogen cyanide and explosives are also highly dangerous. White phosphorus, an allotrope of phosphorus, is toxic, with 100 grams being a lethal dose. White phosphorus usually kills humans within a week of ingestion by attacking the liver. Breathing in phosphorus in its gaseous form can cause an industrial disease called "phossy jaw", which eats away the jawbone. White phosphorus is also highly flammable. Some organophosphorus compounds can fatally block certain enzymes in the human body. Elemental arsenic is toxic, as are many of its inorganic compounds; however some of its organic compounds can promote growth in chickens. The lethal does of arsenic for a typical adult is 200 milligrams, and can cause diarrhea, vomiting, colic, dehydration, and coma. Death from arsenic poisoning typically occurs within a day. Antimony is mildly toxic. Additionally, wine steeped in antimony containers can induce vomiting. When taken in large doses, antimony causes vomiting in a victim, who then appears to recover before dying several days later. Antimony attaches itself to certain enzymes and is difficult to dislodge. Stibine, or SbH3 is far more toxic than pure antimony. Bismuth itself is largely nontoxic, although consuming too much of it can damage the liver. Only one person has ever been reported to have died from bismuth poisoning. However, consumption of soluble bismuth salts can turn a person's gums black. See also ď&#x201A;ˇ

Oxypnictide, including superconductors discovered in 2008.

ď&#x201A;ˇ

Ferropnictide, including oxypnictide superconductors.

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Chapter 5: Chalcogen The chalcogens (/ˈkælkədʒɨnz/) are the chemical elements in group 16 of the periodic table. This group is also known as the oxygen family. It consists of the elements oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and the radioactive element polonium (Po). The synthetic element livermorium (Lv) is predicted to be a chalcogen as well. Often, oxygen is treated separately from the other chalcogens, sometimes even excluded from the scope of the term "chalcogen" altogether, due to its very different chemical behavior from sulfur, selenium, tellurium and polonium. The word "chalcogen" is derived from a combination of the Greek word khalkόs (σαλκόρ) principally meaning copper (the term was also used for bronze/brass, any metal in the poetic sense, ore or coin), and the Latinised Greek word genēs, meaning born or produced. Sulfur has been known since antiquity, and oxygen was recognized as an element in the 18th century. Selenium, tellurium and polonium were discovered in the 19th century, and livermorium in 2000. All of the chalcogens have six valence electrons, leaving them two electrons short of a full outer shell. Their most common oxidation states are −2, +2, +4, and +6. They have relatively low atomic radii, especially the lighter ones. Lighter chalcogens are typically nontoxic in their elemental form, and are often critical to life, while the heavier chalcogens are typically toxic. All of the chalcogens have some role in biological functions, either as a nutrient or a toxin. The lighter chalcogens, such as oxygen and sulfur, are rarely toxic and usually helpful in their pure form. Selenium is an important nutrient but is also commonly toxic. Tellurium often has unpleasant effects (although some organisms can use it), and polonium is always extremely harmful, both in its chemical toxicity and its radioactivity. Sulfur has more than 20 allotropes, oxygen has nine, selenium has at least five, polonium has two, and only one crystal structure of tellurium has so far been discovered. There are numerous organic chalcogen compounds. Not counting oxygen, organic sulfur compounds are generally the most common, followed by organic selenium compounds and organic tellurium compounds. This trend also occurs with chalcogen pnictides and compounds containing chalcogens and carbon group elements. Oxygen is generally extracted from air and sulfur is extracted from oil and natural gas. Selenium and tellurium are produced as byproducts of copper refining. Polonium and livermorium are most available in particle accelerators. The primary use of elemental oxygen is in steelmaking. Sulfur is mostly converted into sulfuric acid, which is heavily used in the

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chemical industry. Selenium's most common application is glassmaking. Tellurium compounds are mostly used in optical disks, electronic devices, and solar cells. Some of polonium's applications are due to its radioactivity. Properties Atomic and physical Chalcogens show similar patterns in electron configuration, especially in the outermost shells, where they all have the same number of valence electrons, resulting in similar trends in chemical behavior: All chalcogens have six valence electrons. All of the solid, stable chalcogens are soft and do not conduct heat well.Electronegativity decreases towards the chalcogens with higher atomic numbers. Density, melting and boiling points, and atomic and ionic radii tend to increase towards the chalcogens with higher atomic numbers. Isotopes Out of the six known chalcogens, one (oxygen) has an atomic number equal to a nuclear magic number, which means that their atomic nuclei tend to have increased stability towards radioactive decay. Oxygen has three stable isotopes, and 14 unstable ones. Sulfur has four stable isotopes, 20 radioactive ones, and one isomer. Selenium has six observationally stable or nearly stable isotopes, 26 radioactive isotopes, and 9 isomers. Tellurium has eight stable or nearly stable isotopes, 31 unstable ones, and 17 isomers. Polonium has 42 isotopes, none of which are stable. It has an additional 28 isomers. In addition to the stable isotopes, some radioactive chalcogen isotopes occur in nature, either because they are decay products, such as 210Po, because they are primordial, such as 82Se, because of cosmic ray spallation, or via nuclear fission of uranium. Livermorium isotopes 290 through 293 have been discovered. The most stable livermorium isotope is 293Lv, which has a halflife of 0.061 seconds. Among the lighter chalcogens (oxygen and sulfur), the most neutron-poor isotopes undergo proton emission, the moderately neutron-poor isotopes undergo electron capture or Î˛+ decay, the moderately neutron-rich isotopes undergo Î˛- decay, and the most neutron rich isotopes undergo neutron emission. The middle chalcogens (selenium and tellurium) have similar decay tendencies as the lighter chalcogens, but their isotopes do not undergo proton emission and some of the most neutron-starved isotopes of tellurium undergo alpha decay. Polonium's isotopes tend to decay with alpha or beta decay. Isotopes with nuclear spins are more common among the chalcogens selenium and tellurium than they are with sulfur.

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Allotropes The four stable chalcogens at STP Oxygen's most common allotrope is diatomic oxygen, or O2, a reactive paramagnetic molecule that is ubiquitous to aerobic organisms and has a blue color in its liquid state. Another allotrope is O3, or ozone, which is three oxygen atoms bonded together in a bent formation. There is also an allotrope called tetraoxygen, or O4, and six allotropes of solid oxygen including "red oxygen", which has the formula O8. Phase diagram for solid oxygen Sulfur has over 20 known allotropes, which is more than any other element except carbon. The most common allotropes are in the form of eight-atom rings, but other molecular allotropes that contain as few as two atoms or as many as 20 are known. Other notable sulfur allotropes include rhombic sulfur and monoclinic sulfur. Rhombic sulfur is the more stable of the two allotropes. Monoclinic sulfur takes the form of long needles and is formed when liquid sulfur is cooled to slightly below its melting point. The atoms in liquid sulfur are generally in the form of long chains, but above 190° Celsius, the chains begin to break down. If liquid sulfur above 190° Celsius is frozen very rapidly, the resulting sulfur is amorphous or "plastic" sulfur. Gaseous sulfur is a mixture of diatomic sulfur (S2) and 8-atom rings. Phase diagram of sulfur showing the relative stabilities of several allotropes Selenium has at least five known allotropes. The gray allotrope, commonly referred to as the "metallic" allotrope, despite not being a metal, is stable and has a hexagonal crystal structure. The gray allotrope of selenium is soft, with a Mohs hardness of 2, and brittle. The four other allotropes of selenium are metastable. These include two monoclinic red allotropes and two amorphous allotropes, one of which is red and one of which is black. The red allotrope converts to the red allotrope in the presence of heat. The gray allotrope of selenium is made from spirals on selenium atoms, while one of the red allotropes is made of stacks of selenium rings (Se8). Tellurium is not known to have any allotropes, although its typical form is hexagonal. Polonium has two allotropes, which are known as α-polonium and β-polonium. α-polonium has a cubic crystal structure and converts the rhombohedral β-polonium at 36 °C. The chalcogens have varying crystal structures. Oxygen's crystal structure is monoclinic, sulfur's is orthorhombic, selenium and tellurium have the hexagonal crystal structure, while polonium has a cubic crystal structure.

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Chemical Oxygen, sulfur, and selenium are nonmetals, and tellurium is a metalloid, meaning that its chemical properties are between those of a metal and those of a nonmetal. It is not certain whether polonium is a metal or a metalloid. Some sources refer to polonium as a metalloid, although it has some metallic properties. Also, some allotropes of selenium display characteristics of a metalloid, even though selenium is usually considered a nonmetal. Even though oxygen is a chalcogen, its chemical properties are different from those of other chalcogens. One reason for this is that the heavier chalcogens have vacant d-orbitals. Oxygen's electronegativity is also much higher than those of the other chalcogens. This makes oxygen's electric polarizability several times lower than those of the other chalcogens. The oxidation number of the most common chalcogen compounds is −2. However the tendency for chalcogens to form compounds in the −2 state decreases towards the heavier chalcogens. Other oxidation numbers, such as −1 in pyrite and peroxide, do occur. The highest formal oxidation number is +6. This oxidation number is found in sulfates, selenates, tellurates, polonates, and their corresponding acids, such as sulfuric acid. There are many acids containing chalcogens, including sulfuric acid, sulfurous acid, selenic acid, and telluric acid. All hydrogen chalcogenides are toxic except for water. Oxygen ions 2−

often come in the forms of oxide ions (O ), peroxide ions (O 2−

2−2

−

), and hydroxide ions (OH ). −3

−4

Sulfur ions generally come in the form of sulfides (S ), sulfites (SO2 ), sulfates (SO2 ), and −3

2−

thiosulfates (S2O2 ). Selenium ions usually come in the form of selenides (Se ) and −4

−4

selenates (SeO2 ). Tellurium ions often come in the form of tellurates (TeO2 ). Molecules containing metal bonded to chalcogens are common as minerals. For example, pyrite (FeS 2) is an iron ore, and the rare mineral calaverite is the ditelluride (Au, Ag)Te2. Water (H2O) is the most familiar chalcogen-containing compound. Oxygen is the most electronegative element except for fluorine, and forms compounds with almost all of the chemical elements, including some of the noble gases. It commonly bonds with many metals and metalloids to form oxides, including iron oxide, titanium oxide, and silicon oxide. Oxygen's most common oxidation state is −2, and the oxidation state −1 is also relatively common. With hydrogen it forms water and hydrogen peroxide. Organic oxygen compounds are ubiquitous in organic chemistry. Sulfur's oxidation states are −2, +2, +4, and +6. Sulfur-containing analogs of oxygen compounds often have the prefix thio-. Sulfur's chemistry is similar to oxygen's, in many ways. One difference is that sulfur double bonds are far weaker than oxygen double bonds, but sulfur single bonds are stronger than oxygen single bonds. Organic sulfur compounds such as thiols have a strong specific smell, and a few are utilized by some organisms.

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Selenium's oxidation states are −2, +4, and +6. Selenium, like most chalcogens, bonds with oxygen. There are some organic selenium compounds, such as selenoproteins. Tellurium's oxidation states are −2, +2, +4, and +6. Tellurium forms the oxides tellurium monoxide, tellurium dioxide, and tellurium trioxide. Polonium's oxidation states are +2 and +4. Although all group 16 elements of the periodic table, including oxygen, can be defined as chalcogens, oxygen and oxides are usually distinguished from

chalcogens

and

chalcogenides. The term chalcogenide is more commonly reserved for sulfides, selenides, and tellurides, rather than for oxides. Binary compounds of the chalcogens are called chalcogenides (rather than chalcides; this breaks the pattern of halogen/halide and pnictogen/pnictide). Except for polonium, the chalcogens are all fairly similar to each other chemically. They all form X2− ions when reacting with electropositive metals. Except for oxygen most chalcogens have common oxidation states of +6, +4, and −2. Sulfide minerals and analogous compounds produce gases upon reaction with oxygen. Compounds Halides Chalcogens also form compounds with halogens. Such compounds are known as chalcogen halides. The majority of simple chalcogen halides are well-known and widely used as chemical reagents. However, more complicated chalcogen halides, such as sulfenyl, sulfonyl, and sulfuryl halides, are less well-known to science. Out of the compounds consisting purely of chalcogens and halogens, there are a total of 13 chalcogen fluorides, nine chalcogen chlorides, eight chalcogen bromides, and six chalcogen iodides that are known. The heavier chalcogen halides often have significant molecular interactions. Sulfur fluorides with low valences are fairly unstable and little is known about their properties. However, sulfur fluorides with high valences, such as sulfur hexafluoride, are stable and wellknown. Sulfur tetrafluoride is also a well-known sulfur fluoride. Certain selenium fluorides, such as selenium difluoride, have been produced in small amounts. The crystal structures of both selenium tetrafluoride and tellurium tetrafluoride are known. Chalcogen chlorides and bromides have also been explored. In particular, selenium dichloride and sulfur dichloride can react to form organic selenium compounds. Dichalcogen dihalides, such as Se2Cl2 also are known to exist. There are also mixed chalcogen-halogen compounds. These include SeSX, with X being chlorine or bromine. Such compounds can form in mixtures of sulfur dichloride and selenium halides. These compounds have been fairly recently structurally

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characterized, as of 2008. In general, diselenium and disulfur chlorides and bromides are useful chemical reagents. Chalcogen halides with attached metal atoms are soluble in organic solutions. One example of such a compound is MoS2Cl3. Unlike selenium chlorides and bromides, selenium iodides have not been isolated, as of 2008, although it is likely that they occur in solution. Diselenium diiodide, however, does occur in equilibrium with selenium atoms and iodine molecules. Some tellurium halides with low valences, such as Te2Cl2 and Te2Br2, form polymers when in the solid state. These tellurium halides can be synthesized by the reduction of pure tellurium with superhydride and reacting the resulting product with tellurium tetrahalides. Ditellurium dihalides tend to get less stable as the halides become lower in atomic number and atomic mass. Tellurium also forms iodides with even fewer iodine atoms than diiodies. These include TeI and Te2I. These compounds have extended structures in the solid state. Halogens and chalcogens can also form halochalcogenate anions. Organic Alcohols, phenols and other similar compounds contain chalcogens. They typically contain oxygen. However, in thiols, selenols and tellurols; sulfur, selenium, and tellurium can replace oxygen in these compounds. Thiols are more well known than selenols or tellurols. Thiols are the most stable chalcogenols and tellurols are the least stable, being unstable in heat or light. Other organic chalcogen compounds include thioethers, selenoethers and telluroethers. Some of these, such as dimethyl sulfide, diethyl sulfide, and dipropyl sulfide are commercially available. Selenoethers are in the form of R2Se or RSeR. Telluroethers such as dimethyl telluride are typically prepared in the same way as thioethers and selenoethers. Organic chalcogen compounds, especially organic sulfur compounds, have the tendency to smell unpleasant. Dimethyl telluride also smells unpleasant, and selenophenol is renowned for its "metaphysical stench". There are also thioketones, selenoketones, and telluroketones. Out of these, thioketones are the most well-studied with 80% of chalcogenoketones papers being about them. Selenoketones make up 16% of such papers and telluroketones make up 4% of them. Thioketones have well-studied non-linear electric and photophysic properties. Selenoketones are less stable than thioketones and telluroketones are less stable than selenoketones. Telluroketones have the highest level of polarity of chalcogenoketones. With metals Elemental chalcogens react with certain lanthanide compounds to form lanthanide clusters rich in chalcogens. Uranium (IV) chalcogenol compounds also exist. There are also transition metal chalcogenols which have potential to serve as catalysts and stabilize nanoparticles. There is a very large number of metal chalcogenides. One of the more recent discoveries in

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this group of compounds is Rb2Te. There are also compounds in which alkali metals and transition metals such as the fourth period transition metals except for copper and zinc. In highly metal-rich metal chalcogenides, such as Lu7Te and Lu8Te have domains of the metal's crystal lattice containing chalcogen atoms. While these compounds do exist, analogous chemicals that contain lanthanum, praseodymium, gadolonium, holmium, terbium, or ytterbium have not been discovered, as of 2008. The boron group metals aluminum, gallium, and indium also form bonds to chalcogens. The Ti3+ ion forms chalcogenide dimers such as TiTl5Se8. Metal chalcogenide dimers also occur as lower tellurides, such as Zr5Te6. With pnictogens Bismuth sulfide, a pnictogen chalcogenide Compounds with chalcogen-phosphorus bonds have been explored for more than 200 years. These compounds include unsophisticated phosphorus chalcogenides as well as large molecules with biological roles and phosphorus-chalcogen compounds with metal clusters. These compounds have numerous applications, including strike-anywhere matches and quantum dots. A total of 130,000 phosphorus-sulfur compounds, 6000 phosphorus-selenium compounds, and 350 phosphorus-tellurium compounds have been discovered. The decrease in the number of chalcogen-phosphorus compounds further down the periodic table is due to diminishing bond strength. Such compounds tend at least one phosphorus atom in the center, surrounded by four chalcogens and side chains. However, some phosphoruschalcogen compounds also contain hydrogen (such as secondary phosphine chalcogenides) or nitrogen (such as dichalcogenoimidodiphosphates). Phosphorus selenides are typically harder to handle that phosphorus sulfides, and compounds in the from PxTey have not been discovered. Chalcogens also bond with other pnictogens, such as arsenic, antimony, and bismuth. Heavier chalcogen pnictides tend to form ribbon-like polymers instead of individual molecules. Chemical formulas of these compounds include Bi2S3 and Sb2Se3. Ternary chalcogen pnictides are also known. Examples of these include P4O6Se and P3SbS3. salts containing chalcogens and pnictogens also exist. Almost all chalcogen pnictide salts are typically in the form of [PxE4x]3-, where E is a chalcogen. Tertiary phosphines can react with chalcogens to form compounds in the form of R3PE, where E is a chalcogen. When E is sulfur, these compounds are relatively stable, but they are less so when E is selenium or tellurium. Similarly, secondary phosphines can react with chalcogens to form secondary phosphine chalcogenides. However, these compounds are in a state of equilibrium with chalcogenophosphinous acid. Secondary phosphine chalcogenides are weak acids. Other Chalcogens form single bonds and double bonds with other carbon group elements than

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carbon, such as silicon, germanium, and tin. Such compounds typically form from a reaction of carbon group halides and chalcogenol salts or chalcogenol bases. Cyclic compounds with chalcogens, carbon group elements, and boron atoms exist, and occur from the reaction of boron dichalcogenates and carbon group metal halides. Compounds in the form of M-E, where M is silicon, germanium, or tin, and E is sulfur, selenium or tellurium have been discovered. These form when carbon group hydrides react or when heavier versions of carbenes react. Sulfur and tellurium can bond with organic compounds containing both silicon and phosphorus. All of the chalcogens form hydrides. In some cases this occurs with chalcogens bonding with two hydrogen atoms. However tellurium hydride and polonium hydride are both volatile and highly labile. Also, oxygen can bond to hydrogen in a 1:1 ratio as in hydrogen peroxide, but this compound is unstable. Chalcogen compounds form a number of interchalcogens. For instance, sulfur forms the toxic sulfur dioxide and sulfur trioxide. Tellurium also forms oxides. There are some chalcogen sulfides as well. These include selenium sulfide, an ingredient in some shampoos. Since 1990, a number of borides with chalcogens bonded to them have been detected. The chalcogens in these compounds are mostly sulfur, although some do contain selenium instead. One such chalcogen boride consists of two molecules of dimethyl sulfide attached to a boron-hydrogen molecule. Other important boron-chalcogen compounds include macropolyhedral systems. Such compounds tend to feature sulfur as the chalcogen. There are also chalcogen borides with two, three, or four chalcogens. Many of these contain sulfur but some, such as Na2B2Se7 contain selenium instead. History Early discoveries Greek fire, an early sulfur-related discovery Sulfur was known in the ancient history and is mentioned in Bible 15 times. Sulfur was known to the ancient Greeks and commonly mined by the ancient Romans. Sulfur was also historically used as a component of Greek fire. In the Middle Ages, sulfur was a key part of alchemical experiments. In the 1700s and 1800s, scientists Joseph Louis Gay-Lussac and Louis-Jacques ThĂŠnard proved sulfur to be a chemical element. Early attempts to discover oxygen from air were hampered by the fact that air was thought of as a single element up to the 17th and 18th centuries. Robert Hooke, Mikhail Lomonosov, Ole Borch, and Pierre Bayden all successfully created oxygen, but did not realize it at the

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time. Oxygen was discovered by Joseph Priestley in 1774 when he focused sunlight on a sample of mercuric oxide and collected the resulting gas. Carl Wilhelm Scheele had also created oxygen in 1771 by the same method, but Scheele did not publish his results until 1777. Tellurium was first discovered in 1783 by Franz Joseph Müller von Reichenstein. He discovered tellurium in a sample of what is now known as calaverite. Müller assumed at first that the sample was pure antimony, but tests he ran on the sample did not agree with this. Muller then guessed that the sample was bismuth sulfide, but tests confirmed that the sample was not that. For some years, Muller pondered the problem. Eventually he realized that the sample was gold bonded with an unknown element. In 1796, Müller sent part of the sample to the German chemist Martin Klaproth, who purified the undiscovered element. Klaproth decided to call the element tellurium after the Latin word for earth. Selenium was discovered in 1817 by Jöns Jacob Berzelius. Berzelius discovered a reddishbrown sediment at a sulfuric acid manufacturing plant. The sample was thought to contain arsenic. Berzelius initially thought that the sediment contained tellurium, but came to realize that the sample also contained a new element, which he named selenium after the Greek word for moon. Periodic table placing Mendeleev's periodic system proposed in 1871 showing oxygen, sulfur, selenium and tellurium part of his group VI Three of the chalcogens (sulfur, selenium, and tellurium) were part of the discovery of periodicity, as they are among a series of triads of elements in the same group that were noted by Johann Wolfgang Döbereiner as having similar properties. Around 1865 John Newlands produced a series of papers where he listed the elements in order of increasing atomic weight and similar physical and chemical properties that recurred at intervals of eight; he likened such periodicity to the octaves of music. His version included a "group b" consisting of oxygen, sulfur, selenium, tellurium, and osmium. Johann Wolfgang Döbereiner was among the first to notice similarities between what are now known as chalcogens. After 1869, Dmitri Mendeleev proposed his periodic table placing oxygen at the top of "group VI" above sulfur, selenium, and tellurium.Chromium, molybdenum, tungsten, and uranium were sometimes included in this group, but they would be later rearranged as part of group VIB; uranium would later be moved to the actinide series. Oxygen, along with sulfur, selenium, tellurium, and later polonium would be grouped in group VIA, until the group's

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name was changed to group 16 in 1988.

Modern discoveries In the late 19th century, Marie Curie and Pierre Curie discovered that a sample of pitchblende was emitting four times as much radioactivity as could be explained by the presence of uranium alone. The Curies gathered several tons of pitchblende and refined it for several months until they had a pure sample of polonium. The discovery officially took place in 1898. Prior to the invention of particle accelerators, the only way to create polonium was to extract it over several months from uranium ore. The first attempt at creating livermorium was from 1976 to 1977 at the LBNL, who bombarded curium-248 with calcium-48, but were not successful. After several failed attempts in 1977, 1998, and 1999 by research groups in Russia, Germany, and the USA, livermorium was created successfully in 2000 at the Joint Institute for Nuclear Research by bombarding curium-248 atoms with calcium-48 atoms. The element was known as ununhexium until it was officially named livermorium in 2012. Etymology In the 19th century, Jons Jacob Berzelius suggested calling the elements in group 16 "amphigens", as the elements in the group formed amphid salts (salts of oxyacids) The term received some use in the early 1800s but is now obsolete. The name chalcogen comes from the Greek words σαλκορ (chalkos, literally "copper"), and γενέρ (genes, born, gender, kindle). It was first used in 1932 by Wilhelm Biltz's group at the University of Hanover, where it was proposed by Werner Fischer. The word "chalcogen" gained popularity in Germany during the 1930s because the term was analogous to "halogen". Although the literal meanings of the Greek words imply that chalcogen means "copper-former", this is misleading because the chalcogens have nothing to do with copper in particular. "Ore-former" has been suggested as a better translation, as the vast majority of metal ores are chalcogenides and the word σαλκορ in ancient Greek was associated with metals and metal-bearing rock in general; copper, and its alloy bronze, was one of the first metals to be used by humans. Oxygen's name comes from the Greek words oxy genes, meaning "acid-forming". Sulfur's name comes from either the Latin word sulfurium or the Sanskrit word sulvere; both of those terms are ancient words for sulfur. Selenium is named after the Greek goddess of the moon, Selene, to match the previously-discovered element tellurium, whose name comes from the Latin word telus, meaning earth. Polonium is named after Marie Curie's country of birth, Poland. Livermorium is named for the Lawrence Livermore National Laboratory.

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Occurrence The four lightest chalcogens (oxygen, sulfur, selenium, and tellurium) are all primordial elements on Earth. Sulfur and oxygen occur as constituent copper ores and selenium and tellurium occur in small traces in such ores.Polonium forms naturally after the decay of other elements, even though it is not primordial. Livermorium does not occur naturally at all. Oxygen makes up 21% of the atmosphere by weight, 89% of water by weight, 46% of the earth's crust by weight, and 65% of the human body. Oxygen also occurs in many minerals, being found in all oxide minerals and hydroxide minerals, and in numerous other mineral groups. Stars of at least eight times the mass of the sun also produce oxygen in their cores via nuclear fusion. Oxygen is the third-most abundant element in the universe, making up 1% of the universe by weight. Sulfur makes up 0.035% of the earth's crust by weight, making it the 17th most abundant element there and makes up 0.25% of the human body. It is a major component of soil. Sulfur makes up 870 parts per million of seawater and about 1 part per billion of the atmosphere. Sulfur can be found in elemental form or in the form of sulfide minerals, sulfate minerals, or sulfosalt minerals. Stars of at least 12 times the mass of the sun produce sulfur in their cores via nuclear fusion. Sulfur is the tenth most abundant element in the universe, making up 500 parts per million of the universe by weight. Selenium makes up 0.05 parts per million of the earth's crust by weight. This makes it the 67th most abundant element in the earth's crust. Selenium makes up on average 5 parts per million of the soils. Seawater contains around 200 parts per trillion of selenium. The atmosphere contains 1 nanogram of selenium per cubic meter. There are mineral groups known as selenates and selenites, but there are not many of minerals in these groups. Selenium is not produced directly by nuclear fusion. Selenium makes up 30 parts per billion of the universe by weight. There are only 5 parts per billion of tellurium in the earth's crust and 15 parts per billion of tellurium in seawater. Tellurium is one of the eight or nine least abundant elements in the earth's crust. There are a few dozen tellurate minerals and telluride minerals, and tellurium occurs in some minerals with gold, such as sylvanite and calaverite. Tellurium makes up 9 parts per billion of the universe by weight. Polonium only occurs in trace amounts on earth, via radioactive decay of uranium and thorium. It is present in uranium ores in concentrations of 100 micrograms per metric ton. Very minute amounts of polonium exist in the soil and thus in most food, and thus in the

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human body. The earth's crust contains less than 1 part per billion of polonium, making it one of the ten rarest metals on earth. Livermorium is always produced artificially in particle accelerators. Even when it is produced, only a small number of atoms at a time are synthesized.

Chalcophile elements Chalcophile elements are those that remain on or close to the surface because they combine readily with chalcogens other than oxygen, forming compounds which do not sink into the core. Chalcophile ("chalcogen-loving") elements in this context are those metals and heavier nonmetals that have a low affinity for oxygen and prefer to bond with the heavier chalcogen sulfur as sulfides. Because sulfide minerals are much denser than the silicate minerals formed by lithophile elements, chalcophile elements separated below the lithophiles at the time of the first crystallisation of the Earth's crust. This has led to their depletion in the Earth's crust relative to their solar abundances, though this depletion has not reached the levels found with siderophile elements. Production Approximately 100 million metric tons of oxygen are produced yearly. Oxygen is most commonly produced by fractional distillation, in which air is cooled to a liquid, then warmed, allowing all the components of air except for oxygen to turn to gases and escape. Fractionally distilling air several times can produce 99.5% pure oxygen. Another method with which oxygen is produced is to send a stream of dry, clean air through a bed of molecular sieves made of zeolite, which absorbs the nitrogen in the air, leaving 90 to 93% pure oxygen. Sulfur recovered from oil refining in Alberta, stockpiled for shipment in North Vancouver, British Columbia Sulfur can be mined in its elemental form, although this method is no longer as popular as it used to be. In 1865 a large deposit of elemental sulfur was discovered in the U.S. states of Louisiana and Texas, but it was difficult to extract at the time. In the 1890s, Herman Frasch came up with the solution of liquefying the sulfur with superheated steam and pumping the sulfur up to the surface. These days sulfur is instead more often extracted from oil, natural gas, and tar. The world production of selenium is around 1500 metric tons per year, out of which roughly 10% is recycled. Japan is the largest producer, producing 800 metric tons of selenium per year. Other large producers include Belgium (300 metric tons per year), the United States

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(over 200 metric tons per year), Sweden (130 metric tons per year), and Russia (100 metric tons per year). Selenium can be extracted from the waste from the process of electrolytically refining copper. Another method of producing selenium is to farm selenium-gathering plants such as milk vetch. This method could produce three kilograms of selenium per acre, but is not commonly practiced. Tellurium is mostly produced as a by-product of the processing of copper. Tellurium can also be refined by electrolytic reduction of sodium telluride. The world production of tellurium is between 150 and 200 metric tons per year. The United States is one of the largest producers of tellurium, producing around 50 metric tons per year. Peru, Japan, and Canada are also large producers of tellurium. Until the creation of nuclear reactors, all polonium had to be extracted from uranium ore. In modern times, most isotopes of polonium are produced by bombarding bismuth with neutrons. Polonium can also be produced by high neutron fluxes in nuclear reactors. Approximately 100 grams of polonium are produced yearly. All the polonium produced for commercial purposes is made in the Ozersk nuclear reactor in Russia. From there, it is taken to Samara, Russia for purification, and from there to St. Petersburg for distribution. The United States is the largest consumer of polonium. All livermorium is produced artificially in particle accelerators. The first successful production of livermorium was achieved by bombarding curium-248 atoms with calcium-48 atoms. As of 2011, roughly 25 atoms of livermorium had been synthesized. Applications Steelmaking is the most important use of oxygen; 55% of all oxygen produced goes to this application. The chemical industry also uses large amounts of oxygen; 25% of all oxygen produced goes to this application. The remaining 20% of oxygen produced is mostly split between medical use, water treatment (as oxygen kills some types of bacteria), rocket fuel (in liquid form), and metal cutting. Most sulfur produced is transformed into sulfur dioxide, which is further transformed into sulfuric acid, a very common industrial chemical. Other common uses include being a key ingredient of gunpowder and Greek fire, and being used to change soil pH. Sulfur is also mixed into rubber to vulcanize it. Sulfur is used in some types of concrete and fireworks. 60% of all sulfuric acid produced is used to generate phosphoric acid. Gunpowder, an application of sulfur Around 40% of all selenium produced goes to glassmaking. 30% of all selenium produced

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goes to metallurgy, including manganese production. 15% of all selenium produced goes to agriculture. Electronics such as photovoltaic materials claim 10% of all selenium produced. Pigments account for 5% of all selenium produced. Historically, machines such as photocopiers and light meters used one-third of all selenium produced, but this application is in steady decline. Tellurium suboxide, a mixture of tellurium and tellurium dioxide, is used in the rewritable data layer of some CD-RW disks and DVD-RW disks. Bismuth telluride is also used in many microelectronic devices, such as photoreceptors. Tellurium is sometimes used as an alternative to sulfur in vulcanized rubber. Cadmium telluride is used as a high-efficiency material in solar panels. Some of polonium's applications relate to the element's radioactivity. For instance, polonium is used as an alpha-particle generator for research. Polonium alloyed with beryllium provides an efficient neutron source. Polonium is also used in nuclear batteries. Most polonium is used in antistatic devices. Livermorium does not have any uses whatsoever due to its extreme rarity and short half-life. Organochalcogen compounds are involved in the semiconductor process. These compounds also feature into ligand chemistry and biochemistry. One application of chalcogens themselves is to manipulate redox couples in supramolar chemistry (chemistry involving noncovalent bond interactions). This application leads on to such applications as crystal packing, assembly of large molecules, and biological recognition of patterns. The secondary bonding interactions of the larger chalcogens, selenium and tellurium, can create organic solventholding acetylene nanotubes. Chalcogen interactions are useful for conformational analysis and stereoelectronic effects, among other things. Chalcogenides with through bonds also have applications. For instance, divalent sulfur can stabilize carbanions, cationic centers, and radical. Chalcogens can confer upon ligands (such as DCTO) properties such as being able to transform Cu (II) to Cu (I). Studying chalcogen interactions gives access to radical cations, which are used in mainstream synthetic chemistry. Metallic redox centers of biological importance are tunable by interactions of ligands containing chalcogens, such as methionine and selenocysteine. Also, chalcogen through-bonds can provide insight about the process of electron transfer. Biological role DNA, an important biological compound containing oxygen Oxygen is needed by almost all organisms for the purpose of generating ATP. It is also a key component of most other biological compounds, such as water, amino acids and DNA. Human blood contains a large amount of oxygen. Human bones contain 28% oxygen.

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Human tissue contains 16% oxygen. A typical 70-kilogram human contais 43 kilograms of oxygen, mostly in the form of water.

All animals need significant amounts of sulfur. Some amino acids, such as cysteine and methionine contain sulfur. Plant roots take up sulfate ions from the soil and reduce it to sulfide ions. Metalloproteins also use sulfur to attach to useful metal atoms in the body and sulfur similarly attaches itself to poisonous metal atoms like cadmium to haul them to the safety of the liver. On average, humans consume 900 milligrams of sulfur each day. Sulfur compounds, such as those found in skunk spray often have strong odors. All animals and some plants need trace amounts of selenium, but only for some specialized enzymes. Humans consume on average between 6 and 200 micrograms of selenium per day. Mushrooms and brazil nuts are especially noted for their high selenium content. Selenium in foods is most commonly found in the form of amino acids such as selenocysteine and selenomethionine. Selenium can protect against heavy metal poisoning. Tellurium is not known to be needed for animal life, although a few fungi can incorporate it in compounds in place of selenium. Microorganisms also absorb tellurium and emit dimethyl telluride. Most tellurium in the blood stream is excreted slowly in urine, but some is converted to dimethyl telluride and released through the lungs. On average, humans ingest about 600 micrograms of tellurium daily. Plants can take up some tellurium from the soil. Onions and garlic have been found to contain as much as 300 parts per million of tellurium in dry weight. Polonium has no biological role, and is highly toxic on account of being radioactive. Fire diamond for the chalcogen selenium Oxygen is generally nontoxic, but oxygen toxicity has been reported when it is used in high concentrations. In both elemental gaseous form and as a component of water, it is vital to almost all life on earth. Despite this, liquid oxygen is highly dangerous. Even gaseous oxygen is dangerous in excess. For instance, sports divers have occasionally drowned from convulsions caused by breathing pure oxygen at a depth of more than 10 meters (33 feet) underwater. Oxygen is also toxic to some bacteria. Ozone, an allotrope of oxygen, is toxic to most life. It can cause lesions in the respiratory tract. Sulfur is generally nontoxic and is even a vital nutrient for humans. However, in its elemental form it can cause redness in the eyes and skin, a burning sensation and a cough if inhaled, a burning sensation and diarrhea if ingested, and can irritate the mucous membranes. An excess of sulfur can be toxic for cows because microbes in the rumens of cows produce

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toxic hydrogen sulfide upon reaction with sulfur. Many sulfur compounds, such as hydrogen sulfide (H2S) and sulfur dioxide (SO2) are highly toxic. Selenium is a trace nutrient required by humans on the order of tens or hundreds of micrograms per day. A dose of over 450 micrograms can be toxic, resulting in bad breath and body odor. Extended, low-level exposure, which can occur at some industries, results in weight loss, anemia, and dermatitis. In many cases of selenium poisoning, selenous acid is formed in the body.Hydrogen selenide (H2Se) is highly toxic. Tellurium is not generally highly toxic, but can produce unpleasant side effects. As little as 10 micrograms of tellurium per cubic meter of air can cause notoriously unpleasant breath, described as smelling like rotten garlic. Acute tellurium poisoning can cause vomiting, gut inflammation, internal bleeding, and respiratory failure. Extended, low-level exposure to tellurium causes tiredness and indigestion. Sodium tellurite (Na2TeO3) is lethal in amounts of around 2 grams. Polonium is dangerous both as an alpha particle emitter and because it is chemically toxic. If ingested, polonium-210 is a billion times as toxic as hydrogen cyanide by weight; it has been used as a murder weapon in the past, most famously to kill Alexander Litvinenko. Polonium poisoning can cause nausea, vomiting, anorexia, and lymphopenia. It can also damage hair follicles and white blood cells. Polonium-210 is only dangerous if ingested or inhaled because its alpha particle emissions cannot penetrate human skin. Polonium-209 is also toxic, and can cause leukemia.

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Chapter 6: Halogen

Halogens

chalcogens ←

→ noble gases

IUPAC group number

Name by element

fluorine group

Trivial name

halogens

CAS group number (US)

VIIA

old IUPAC number (European)VIIB

↓ Period

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Fluorine

(F)

9 Halogen 3

Chlorine

(Cl)

17 Halogen 4

Bromine

(Br)

35 Halogen

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Iodine

(I)

53 Halogen 6 Astatine

(At)

85 Halogen

Legend primordial element element from decay Atomic number color: black=solid, green=liquid, red=gas

v t e The halogens or halogen elements (/ˈhælɵdʒɨn/) are a group in the periodic table consisting of five chemically related elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The artificially created element 117 (ununseptium) may also be a halogen. In the modern IUPAC nomenclature, this group is known as group 17. The group of halogens is the only periodic table group that contains elements in all three familiar states of matter at standard temperature and pressure. All of the halogens form acids when bonded to hydrogen. Most halogens are typically produced from minerals or salts. The middle halogens, that is, chlorine, bromine and iodine, are often used as disinfectants. The halogens are also all toxic.

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History The fluorine mineral fluorospar was known as early as 1529. Early chemists realized that fluorine compounds contain an undiscovered element, but were unable to isolate it. In 1869, George Gore, an English chemist, ran a current of electricity through hydrofluoric acid and discovered fluorine, but he was unable to prove his results at the time. In 1886, Henri Moissan, a chemist in Paris, performed electrolysis on potassium bifluoride dissolved in waterless hydrofluoric acid, and successfully produced fluorine. Hydrochloric acid was known to alchemists and early chemists. However, elemental chlorine was not produced until 1774, when Carl Wilhelm Scheele heated hydrochloric acid with manganese dioxide. Scheele called the element "dephlogisticated muriatic acid", which is how chlorine was known for 33 years. In 1807, Humphry Davy investigated chlorine and discovered that it is an actual element. Chlorine was used as a poison gas during World War I. Bromine was discovered in the 1820s by Antoine-Jérôme Balard. Balard discovered bromine by passing chlorine gas through a sample of brine. He originally proposed the name muride for the new element, but the French Academy changed the element's name to bromine. Iodine was discovered by Bernard Courtois, who was using seaweed ash as part of a process for saltpeter manufacture. Courtois typically boiled the seaweed ash with water to generate potassium chloride. However, in 1811, Courtois added sulfuric acid to his process, and found that his process produced purple fumes that condensed into black crystals of iodine. Suspecting that he had discovered a new element, Courtois sent his sample to other chemists for investigation. Iodine was proven to be a new element by Joseph Gay-Lussac. In 1931, Fred Allison claimed to have discovered element 85 with a magneto-optical machine, and named the element Alabamine, but was mistaken. In 1937, Jajendralal De claimed to have discovered element 85 in minerals, and called the element dakine, but he was also mistaken. An attempt at discovering element 85 in 1939 by Horia Hulublei and Yvette Cauchois via spectroscopy was also unsuccessful, as was an attempt in the same year by Walter Minder, who discovered an iodine-like element resulting from beta decay of radium. Element 85, now named astatine, was produced successfully in 1940 by Dale R. Corson, K.R. Mackenzie, and Emilio G. Segrè, who bombarded bismuth with alpha particles.

Etymology In 1842, the Swedish chemist Baron Jöns Jakob Berzelius proposed the term "halogen" – ἅλρ (háls), "salt" or "sea", and γεν- (gen-), from γίγνομαι (gígnomai), "come to be" – for the four elements (fluorine, chlorine, bromine, and iodine) that produce a sea-salt-like substance

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when they form a compound with a metal. The word "halogen" had actually first been proposed in 1811 by Johann Salomo Christoph Schweigger as a name for the newly discovered element chlorine, but Davy's proposed term for this element eventually won out, and Schweigger's term was kept at Berzelius' suggestion as the term for the element group that contains chlorine. Fluorine's name comes from the Latin word fluere, meaning "to flow". Chlorine's name comes from the Greek word chloros, meaning "greenish-yellow". Bromine's name comes from the Greek word bromos, meaning "stench". Iodine's name comes from the Greek word iodes, meaning "violet". Astatine's name comes from the Greek word astatos, meaning "unstable".

Characteristics

Chemical The halogens show trends in chemical bond energy moving from top to bottom of the periodic table column with fluorine deviating slightly. (It follows trend in having the highest bond energy in compounds with other atoms, but it has very weak bonds within the diatomic F2 element molecules.) Halogen bond energies (kJ/mol) X

BX3 AlX3 CX4

159 574 645 582

456

Cl 243 428 444 427

327

Br 193 363 368 360

272

239

151 294 272 285

Halogens are highly reactive, and as such can be harmful or lethal to biological organisms in sufficient quantities. This high reactivity is due to the high electronegativity of the atoms due to their high effective nuclear charge. They can gain an electron by reacting with atoms of other elements. Fluorine is one of the most reactive elements in existence, attacking otherwise-inert materials such as glass, and forming compounds with the heavier noble gases. It is a corrosive and highly toxic gas. The reactivity of fluorine is such that, if used or stored in laboratory glassware, it can react with glass in the presence of small amounts of water to form silicon tetrafluoride (SiF4). Thus, fluorine must be handled with substances such as Teflon (which is itself an organofluorine compound), extremely dry glass, or metals such as copper or steel, which form a protective layer of fluoride on their surface.

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The high reactivity of fluorine means that, once it does react with something, it bonds with it so strongly that the resulting molecule is very inert and non-reactive to anything else. For example, Teflon is fluorine bonded with carbon.

Molecules Diatomic halogen molecules The halogens form homonuclear diatomic molecules (not proven for astatine). As such they form part of the group known as "elemental gases".

halogen molecule structure

model

d(Xâ&#x2C6;&#x2019;X) / pm d(Xâ&#x2C6;&#x2019;X) / pm (gas phase) (solid phase)

fluorine F2

143

149

chlorine Cl2

199

198

bromine Br2

228

227

iodine

266

272

astatine At2 The elements become less reactive and have higher melting points as the atomic number increases.

Compounds Hydrogen halides All of the halogens have been observed to react with hydrogen to form hydrogen halides. For

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fluorine, chlorine, and bromine, this reaction is in the form of: H2 + X2 → 2HX However, hydrogen iodide and hydrogen astatide can split back into their constituent elements. The hydrogen-halogen reactions get gradually less reactive toward the heavier halogens. A fluorine-hydrogen reaction is explosive even when it is dark and cold. A chlorine-hydrogen reaction is also explosive, but only in the presence of light and heat. A bromine-hydrogen reaction is even less explosive; it is explosive only when exposed to flames. Iodine and astatine only partially react with hydrogen, forming equilibria. All halogens form binary compounds with hydrogen known as the hydrogen halides: hydrogen fluoride (HF), hydrogen chloride (HCl), hydrogen bromide (HBr), hydrogen iodide (HI), and hydrogen astatide (HAt). All of these compounds form acids when mixed with water. Hydrogen fluoride is the only hydrogen halide that forms hydrogen bonds. Hydrochloric acid, hydrobromic acid, hydroiodic acid, and hydroastatic acid are all strong acids, but hydrofluoric acid is a weak acid. All of the hydrogen halides are irritants. Hydrogen fluoride and hydrogen chloride are highly acidic. Hydrogen fluoride is used as an industrial chemical, and is highly toxic, causing pulmonary edema and damaging cells. Hydrogen chloride is also a dangerous chemical. Breathing in gas with more than fifty parts per million of hydrogen chloride can cause death in humans. Hydrogen bromide is even more toxic and irritating than hydrogen chloride. Breathing in gas with more than thirty parts per million of hydrogen bromide can be lethal to humans. Hydrogen iodide, like other hydrogen halides, is toxic.

Metal halides Main article: Metal halides All the halogens are known to react with sodium to form sodium fluoride, sodium chloride, sodium bromide, sodium iodide, and sodium astatide. Heated sodium's reaction with halogens produces bright-orange flames. Sodium's reaction with chlorine is in the form of: 2Na + Cl2 → 2NaCl Iron reacts with fluorine, chlorine, and bromine to form Iron(III) halides. These reactions are in the form of: 2Fe + 3X2 → 2FeX3

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However, when iron reacts with iodine, it forms only iron(II) iodide. Iron wool can react rapidly with fluorine to form the white compound iron(III) fluoride even in cold temperatures. When chlorine comes into contact with heated iron, they react to form the black iron (III) chloride. However, if the reaction conditions are moist, this reaction will instead result in a reddish-brown product. Iron can also react with bromine to form iron(III) bromide. This compound is reddish-brown in dry conditions. Iron's reaction with bromine is less reactive than its reaction with fluorine or chlorine. Hot iron can also react with iodine, but it forms iron (II) iodide. This compound may be gray, but the reaction is always contaminated with excess iodine, so it is not known for sure. Iron's reaction with iodine is less vigorous than its reaction with the lighter halogens.

Interhalogen compounds Interhalogen compounds are in the form of XYn where X and Y are halogens and n is one, three, five, or seven. Interhalogen compounds contain at most two different halogens. Large interhalogens, such as ClF 3 can be produced by a reaction of a pure halogen with a smaller interhalogen such as ClF. All interhalogens except IF 7 can be produced by directly combining pure halogens in various conditions. Interhalogens are typically more reactive than all diatomic halogen molecules except F 2 because interhalogen bonds are weaker. However, the chemical properties of interhalogens are still roughly the same as those of diatomic halogens. Many interhalogens consist of one or more atoms of fluorine bonding to a heavier halogen. Chlorine can bond with up to 3 fluorine atoms, bromine can bond with up to five fluorine atoms, and iodine can bond with up to seven fluorine atoms. Most interhalogen compounds are covalent gases. However, there are some interhalogens that are liquids, such as BrF 3, and many iodine-containing interhalogens are solids.

Organohalogen compounds Many synthetic organic compounds such as plastic polymers, and a few natural ones, contain halogen atoms; these are known as halogenated compounds or organic halides. Chlorine is by far the most abundant of the halogens, and the only one needed in relatively large amounts (as chloride ions) by humans. For example, chloride ions play a key role in brain function by mediating the action of the inhibitory transmitter GABA and are also used by the body to produce stomach acid. Iodine is needed in trace amounts for the production of thyroid hormones such as thyroxine. On the other hand, neither fluorine nor bromine is believed to be essential for humans. Organohalogens are also synthesized through the nucleophilic abstraction reaction.

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Polyhalogenated compounds Polyhalogenated compounds are industrially created compounds substituted with multiple halogens. Many of them are very toxic and bioaccumulate in humans, and have a very wide application range. They include the much-maligned PCBs, PBDEs, and perfluorinated compounds (PFCs), as well as numerous other compounds.

Reactions Reactions with water Fluorine reacts vigorously with water to produce oxygen (O 2) and hydrogen fluoride (HF): 2 F2(g) + 2 H2O(l) → O2(g) + 4 HF(aq) Chlorine has maximum solubility of ca. 7.1 g Cl2 per kg of water at ambient temperature (21 °C). Dissolved chlorine reacts to form hydrochloric acid (HCl) and hypochlorous acid, a solution that can be used as a disinfectant or bleach: Cl2(g) + H2O(l) → HCl(aq) + HClO(aq) Bromine has a solubility of 3.41 g per 100 g of water, but it slowly reacts to form hydrogen bromide (HBr) and hypobromous acid (HBrO): Br2(g) + H2O(l) → HBr(aq) + HBrO(aq) Iodine, however, is minimally soluble in water (0.03 g/100 g water at 20 °C) and does not react with it. However, iodine will form an aqueous solution in the presence of iodide ion, such as by addition of potassium iodide (KI), because the triiodide ion is formed.

Physical and atomic The table below is a summary of the key physical and atomic properties of the halogens. Data marked with question marks are either uncertain or are estimations partially based on periodic trends rather than observations. Halogen Standard

Melting Melting Boiling Boiling Density Electronegativity First 3

Covalent

atomic

point

(g/cm at (Pauling)

ionization radius

weight

(K)

(°C)

(K)

(°C)

25 °C)

energy

(u) Fluorine 18.9984032(5) 53.53

(pm) −1

(kJ·mol ) −219.62 85.03

−188.12 0.0017

3.98

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Chlorine [35.446;

171.6

−101.5 239.11 −34.04 0.0032

3.16

1251.2

265.8

−7.3

35.457] Bromine 79.904(1) Iodine

126.90447(3) 386.85 113.7

Astatine [210]

575

302

332.0

58.8

3.1028

2.96

1139.9

114

457.4

184.3

4.933

2.66

1008.4

133

? 610

? 337

? 887.7

6.2– 2.2

6.5

Isotopes Fluorine has one stable and naturally occurring isotope, fluorine-19. However, there are trace amounts in nature of the radioactive isotope fluorine-23, which occurs via cluster decay of protactinium-231. A total of eighteen isotopes of fluorine have been discovered, with atomic masses ranging from 14 to 31. Chlorine has two stable and naturally occurring isotopes, chlorine-35 and chlorine-37. However, there are trace amounts in nature of the isotope chlorine-36, which occurs via spallation of argon-36. A total of 24 isotopes of chlorine have been discovered, with atomic masses ranging from 28 to 51. There are two stable and naturally occurring isotopes of bromine, bromine-79 and bromine81. A total of 32 isotopes of bromine have been discovered, with atomic masses ranging 67 to 98. There is one stable and naturally occurring isotope of iodine, iodine-127. However, there are trace amounts in nature of the radioactive isotope iodine-129, which occurs via spallation and from the radioactive decay of uranium in ores. Several other radioactive isotopes of iodine have also been created naturally via the decay of uranium. A total of 38 isotopes of iodine have been discovered, with atomic masses ranging from 108 to 145. There are no stable isotopes of astatine. However, there are three naturally occurring radioactive isotopes of astatine produced via radioactive decay of uranium, neptunium, and plutonium. These isotopes are astatine-215, astatine-217, and astatine-219. A total of 31 isotopes of astatine have been discovered, with atomic masses ranging from 193 to 223.

Production Approximately six million metric tons of the fluorine mineral fluorite are produced each year. Four hundred-thousand metric tons of hydrofluoric acid are made each year. Fluorine gas is made from hydrofluoric acid produced as a by-product of phosphoric acid manufacture. Approximately 15,000 metric tons of fluorine gas are made per year. The mineral halite is the mineral that is most commonly mined for chlorine, but the minerals

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carnallite and sylvite are also mined for chlorine. Forty million metric tons of chlorine are produced each year by the electrolysis of brine. Approximately 450,000 metric tons of bromine are produced each year. Fifty percent of all bromine produced is produced in the United States, 35% in Israel, and most of the remainder in China. Historically, bromine was produced by adding sulfuric acid and bleaching powder to natural brine. However, in modern times, bromine is produced by electrolysis, a method invented by Herbert Dow. It is also possible to produce bromine by passing chlorine through seawater and then passing air through the seawater. In 2003, 22,000 metric tons of iodine were produced. Chile produces 40% of all iodine produced, Japan produces 30%, and smaller amonts are produced in Russia and the United States. Until the 1950s, iodine was extracted from kelp. However, in modern times, iodine is produced in other ways. One way that iodine is produced is by mixing sulfur dioxide with nitrate ores, which contain some iodates. Iodine is also extracted from natural gas fields. Even though astatine is naturally occurring, it is usually produced by bombarding bismuth with alpha particles.

From left to right: chlorine, bromine, and iodine at room temperature. Chlorine is a gas, bromine is a liquid, and iodine is a solid. Fluorine could not be included in the image due to its high reactivity.

Applications Both chlorine and bromine are used as disinfectants for drinking water, swimming pools, fresh wounds, spas, dishes, and surfaces. They kill bacteria and other potentially harmful microorganisms through a process known as sterilization. Their reactivity is also put to use in bleaching. Sodium hypochlorite, which is produced from chlorine, is the active ingredient of most fabric bleaches, and chlorine-derived bleaches are used in the production of some paper products. Chlorine also reacts with sodium to create sodium chloride, which is another name for table salt.

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In drug discovery, the incorporation of halogen atoms into a lead drug candidate results in analogues that are usually more lipophilic and less water-soluble. As a consequence, halogen atoms are used to improve penetration through lipid membranes and tissues. It follows that there is a tendency for some halogenated drugs to accumulate in adipose tissue. The chemical reactivity of halogen atoms depends on both their point of attachment to the lead and the nature of the halogen. Aromatic halogen groups are far less reactive than aliphatic halogen groups, which can exhibit considerable chemical reactivity. For aliphatic carbon-halogen bonds, the C-F bond is the strongest and usually less chemically reactive than aliphatic C-H bonds. The other aliphatic-halogen bonds are weaker, their reactivity increasing down the periodic table. They are usually more chemically reactive than aliphatic C-H bonds. As a consequence, the most common halogen substitutions are the less reactive aromatic fluorine and chlorine groups.

Biological role Fluoride anions are found in ivory, bones, teeth, blood, eggs, urine, and hair of organisms. Fluoride anions in very small amounts are essential for humans. There are 0.5 milligrams per liter of fluorine in human blood. Human bones contain 0.2 to 1.2% fluorine. Human tissue contains approximately 50 parts per billion of fluorine. A typical 70-kilogram human contains 3 to 6 grams of fluorine. Chloride anions are essential to a large number of species, humans included. The concentration of chlorine in the dry weight of cereals is 10 to 20 parts per million, while in potatoes the concentration of chloride is 0.5%. Plant growth is adversely affected by chloride levels in the soil falling below 2 parts per million. Human blood contains an average of 0.3% chlorine. Human bone contains typically contains 900 parts per million of chlorine. Human tissue contains approximately 0.2 to 0.5% chlorine. There is a total of 95 grams of chlorine in a typical 70-kilogram human. Some bromine in the form of the bromide anion is present in all organisms. A biological role for bromine in humans has not been proven, but some organisms contain organobromine compounds. Humans typically consume 1 to 20 milligrams of bromine per day. There are typically 5 parts per million of bromine in human blood, 7 parts per million of bromine in human bones, and 7 parts per million of bromine in human tissue. A typical 70-kilogram human contains 260 milligrams of bromine. Humans typically consume less than 100 micrograms of iodine per day. Iodine deficiency can cause intellectual disability. Organoiodine compounds occur in humans in some of the glands, especially the thyroid gland, as well as the stomach, epidermis, and immune system.

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Foods containing iodine include cod, oysters, shrimp, herring, lobsters, sunflower seeds, seaweed, and mushrooms. However, iodine is not known to have a biological role in plants. There are typically 0.06 milligrams per liter of iodine in human blood, 300 parts per billion of iodine in human bones, and 50 to 700 parts per billion of iodine in human tissue. There are 10 to 20 milligrams of iodine in a typical 70-kilogram human. Astatine has no biological role.

Toxicity The halogens tend to decrease in toxicity towards the heavier halogens. Fluorine gas is extremely toxic; breathing fluorine gas at a concentration of 0.1% for several minutes is lethal. Hydrofluoric acid is also toxic, being able to penetrate skin and cause highly painful burns. In addition, fluoride anions are toxic, but not as toxic as pure fluorine. Fluoride can be lethal in amounts of 5 to 10 grams. Prolonged consumption of fluoride above concentrations of 1.5 mg/L is associated with a risk of dental fluorosis, an aesthetic condition of the teeth. At concentrations above 4 mg/L, there is an increased risk of developing skeletal fluorosis, a condition in which bone fractures become more common due to the hardening of bones. Current recommended levels in water fluoridation, a way to prevent dental caries, range from 0.7-1.2 mg/L to avoid the detrimental effects of fluoride while at the same time reaping the benefits. People with levels between normal levels and those required for skeletal fluorosis tend to have symptoms similar to arthritis. Chlorine gas is highly toxic. Breathing in chlorine at a concentration of 3 parts per million can rapidly cause a toxic reaction. Breathing in chlorine at a concentration of 50 parts per million is highly dangerous. Breathing in chlorine at a concentration of 500 parts per million for a few minutes is lethal. Breathing in chlorine gas is highly painful.Hydrochloric acid is a dangerous chemical. Pure bromine is somewhat toxic, but less toxic than fluorine and chlorine. One hundred milligrams of bromine are lethal. Bromide anions are also toxic, but less so than bromine. Bromide has a lethal dose of 30 grams. Iodine is somewhat toxic, being able to irritate the lungs and eyes, with a safety limit of 1 milligram per cubic meter. When taken orally, 3 grams of iodine can be lethal. Iodide anions are mostly nontoxic, but these can also be deadly if ingested in large amounts. Astatine is very radioactive and thus highly dangerous.

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See also 

Pseudohalogen



Halogen bond



Halogen lamp



Interhalogen

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Chapter 7: Noble gas The noble gases make a group of chemical elements with similar properties: under standard conditions, they are all odorless, colorless, monatomic gases with very low chemical reactivity. The six noble gases that occur naturally are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive radon (Rn). For the first six periods of the periodic table, the noble gases are exactly the members of group 18 of the periodic table. It is possible that due to relativistic effects, the group 14 element flerovium exhibits some noble-gas-like properties, instead of the group 18 element ununoctium. The properties of the noble gases can be well explained by modern theories of atomic structure: their outer shell of valence electrons is considered to be "full", giving them little tendency to participate in chemical reactions, and it has been possible to prepare only a few hundred noble gas compounds. The melting and boiling points for a given noble gas are close together, differing by less than 10 째C (18 째F); that is, they are liquids over only a small temperature range. Neon, argon, krypton, and xenon are obtained from air in an air separation unit using the methods of liquefaction of gases and fractional distillation. Helium is sourced from natural gas fields which have high concentrations of helium in the natural gas, using cryogenic gas separation techniques, and radon is usually isolated from the radioactive decay of dissolved radium compounds. Noble gases have several important applications in industries such as lighting, welding, and space exploration. A helium-oxygen breathing gas is often used by deep-sea divers at depths of seawater over 55 m (180 ft) to keep the diver from experiencing oxygen toxemia, the lethal effect of high-pressure oxygen, and nitrogen narcosis, the distracting narcotic effect of the nitrogen in air beyond this partial-pressure threshold. After the risks caused by the flammability of hydrogen became apparent, it was replaced with helium in blimps and balloons. History Noble gas is translated from the German noun Edelgas, first used in 1898 by Hugo Erdmann to indicate their extremely low level of reactivity. The name makes an analogy to the term "noble metals", which also have low reactivity. The noble gases have also been referred to as inert gases, but this label is deprecated as many noble gas compounds are now known.Rare gases is another term that was used, but this is also inaccurate because argon forms a fairly considerable part (0.94% by volume, 1.3% by mass) of the Earth's atmosphere.

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Helium was first detected in the Sun due to its characteristic spectral lines. Pierre Janssen and Joseph Norman Lockyer discovered a new element on August 18, 1868 while looking at the chromosphere of the Sun, and named it helium after the Greek word for the Sun, ήλιορ (ílios or helios). No chemical analysis was possible at the time, but helium was later found to be a noble gas. Before them, in 1784, the English chemist and physicist Henry Cavendish had discovered that air contains a small proportion of a substance less reactive than nitrogen. A century later, in 1895, Lord Rayleigh discovered that samples of nitrogen from the air were of a different density than nitrogen resulting from chemical reactions. Along with Scottish scientist William Ramsay at University College, London, Lord Rayleigh theorized that the nitrogen extracted from air was mixed with another gas, leading to an experiment that successfully isolated a new element, argon, from the Greek word απγόρ (argós, "inactive"). With this discovery, they realized an entire class of gases was missing from the periodic table. During his search for argon, Ramsay also managed to isolate helium for the first time while heating cleveite, a mineral. In 1902, having accepted the evidence for the elements helium and argon, Dmitri Mendeleev included these noble gases as group 0 in his arrangement of the elements, which would later become the periodic table. Ramsay continued to search for these gases using the method of fractional distillation to separate liquid air into several components. In 1898, he discovered the elements krypton, neon, and xenon, and named them after the Greek words κπςπτόρ (kryptós, "hidden"), νέορ (néos, "new"), and ξένορ (xénos, "stranger"), respectively. Radon was first identified in 1898 by Friedrich Ernst Dorn, and was named radium emanation, but was not considered a noble gas until 1904 when its characteristics were found to be similar to those of other noble gases. Rayleigh and Ramsay received the 1904 Nobel Prizes in Physics and in Chemistry, respectively, for their discovery of the noble gases; in the words of J. E. Cederblom, then president of the Royal Swedish Academy of Sciences, "the discovery of an entirely new group of elements, of which no single representative had been known with any certainty, is something utterly unique in the history of chemistry, being intrinsically an advance in science of peculiar significance". The discovery of the noble gases aided in the development of a general understanding of atomic structure. In 1895, French chemist Henri Moissan attempted to form a reaction between fluorine, the most electronegative element, and argon, one of the noble gases, but failed. Scientists were unable to prepare compounds of argon until the end of the 20th

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century, but these attempts helped to develop new theories of atomic structure. Learning from these experiments, Danish physicist Niels Bohr proposed in 1913 that the electrons in atoms are arranged in shells surrounding the nucleus, and that for all noble gases except helium the outermost shell always contains eight electrons. In 1916, Gilbert N. Lewis formulated the octet rule, which concluded an octet of electrons in the outer shell was the most stable arrangement for any atom; this arrangement caused them to be unreactive with other elements since they did not require any more electrons to complete their outer shell. In 1962 Neil Bartlett discovered the first chemical compound of a noble gas, xenon hexafluoroplatinate. Compounds of other noble gases were discovered soon after: in 1962 for

radon,

radon

difluoride,

and

1963

for

krypton,

krypton

difluoride

(KrF

2). The first stable compound of argon was reported in 2000 when argon fluorohydride (HArF) was formed at a temperature of 40 K (−233.2 °C; −387.7 °F). In December 1998, scientists at the Joint Institute for Nuclear Research working in Dubna, Russia bombarded plutonium (Pu) with calcium (Ca) to produce a single atom of element 114,flerovium (Fl). Preliminary chemistry experiments have indicated this element may be the first superheavy element to show abnormal noble-gas-like properties, even though it is a member of group 14 on the periodic table. In October 2006, scientists from the Joint Institute for Nuclear Research and Lawrence Livermore National Laboratory successfully created synthetically ununoctium (Uuo), the seventh element in group 18, by bombarding californium (Cf) with calcium (Ca). Physical and atomic properties Property

Helium Neon

Density (g/dm³)

0.1786 0.9002 1.7818 3.708

5.851 9.97

Boiling point (K)

4.4

27.3

87.4

121.5

166.6 211.5

Melting point (K)

0.95

24.7

83.6

115.8

161.7 202.2

Enthalpy of vaporization (kJ/mol)

0.08

1.74

6.52

9.05

12.65 18.1

Solubility in water at 20 °C (cm /kg)

8.61

10.5

33.6

59.4

108.1 230

Atomic number

Atomic radius (calculated) (pm)

108

120

Ionization energy (kJ/mol)

2372

2080

1520

1351

1170

1037

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Krypton Xenon Radon

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Allen electronegativity

4.16

4.79

3.24

2.97

2.58

The noble gases have weak interatomic force, and consequently have very low melting and boiling points. They are all monatomic gases under standard conditions, including the elements with larger atomic masses than many normally solid elements. Helium has several unique qualities when compared with other elements: its boiling and melting points are lower than those of any other known substance; it is the only element known to exhibit superfluidity; it is the only element that cannot be solidified by cooling under standard conditions—a pressure of 25 standard atmospheres (2,500 kPa; 370 psi) must be applied at a temperature of 0.95 K (−272.200 °C; −457.960 °F) to convert it to a solid. The noble gases up to xenon have multiple stable isotopes. Radon has no stable isotopes; its longest-lived isotope,

222

Rn, has a half-life of 3.8 days and decays to form helium and polonium, which

ultimately decays to lead.

This is a plot of ionization potential versus atomic number. The noble gases, which are labeled, have the largest ionization potential for each period. The noble gas atoms, like atoms in most groups, increase steadily in atomic radius from one period to the next due to the increasing number of electrons. The size of the atom is related to several properties. For example, the ionization potential decreases with an increasing radius because the valence electrons in the larger noble gases are farther away from the nucleus and are therefore not held as tightly together by the atom. Noble gases have the largest ionization potential among the elements of each period, which reflects the stability of their electron configuration and is related to their relative lack of chemical reactivity. Some of the heavier noble gases, however, have ionization potentials small enough to be comparable to those of other elements and molecules. It was the insight that xenon has an ionization potential similar to that of the oxygen molecule that led Bartlett to attempt oxidizing xenon using platinum hexafluoride, an oxidizing agent known to be strong enough to react with oxygen. Noble gases cannot accept an electron to form stable anions; that is, they have a negative electron affinity.

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2.60

The macroscopic physical properties of the noble gases are dominated by the weak van der Waals forces between the atoms. The attractive force increases with the size of the atom as a result of the increase in polarizability and the decrease in ionization potential. This results in systematic group trends: as one goes down group 18, the atomic radius, and with it the interatomic forces, increases, resulting in an increasing melting point, boiling point, enthalpy of vaporization, and solubility. The increase in density is due to the increase in atomic mass. The noble gases are nearly ideal gases under standard conditions, but their deviations from the ideal gas law provided important clues for the study of intermolecular interactions. The Lennard-Jones potential, often used to model intermolecular interactions, was deduced in 1924 by John Lennard-Jones from experimental data on argon before the development of quantum mechanics provided the tools for understanding intermolecular forces from first principles. The theoretical analysis of these interactions became tractable because the noble gases are monatomic and the atoms spherical, which means that the interaction between the atoms is independent of direction, or isotropic. Chemical properties

Neon, like all noble gases, has a full valence shell. Noble gases have eight electrons in the outermost shell, except in the case of helium, which has two. The noble gases are colorless, odorless, tasteless, and nonflammable under standard conditions. They were once labeled group 0 in the periodic table because it was believed they had a valence of zero, meaning their atoms cannot combine with those of other elements to form compounds. However, it was later discovered some do indeed form compounds, causing this label to fall into disuse. Like other groups, the members of this family show patterns in its electron configuration, especially the outermost shells resulting in trends in chemical behavior:

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Element

No. of electrons/shell

helium

10 neon

2, 8

18 argon

2, 8, 8

36 krypton

2, 8, 18, 8

54 xenon

2, 8, 18, 18, 8

86 radon

2, 8, 18, 32, 18, 8

The noble gases have full valence electron shells. Valence electrons are the outermost electrons of an atom and are normally the only electrons that participate in chemical bonding. Atoms with full valence electron shells are extremely stable and therefore do not tend to form chemical bonds and have little tendency to gain or lose electrons. However, heavier noble gases such as radon are held less firmly together by electromagnetic force than lighter noble gases such as helium, making it easier to remove outer electrons from heavy noble gases. Noble gas notation As a result of a full shell, the noble gases can be used in conjunction with the electron configuration notation to form the noble gas notation. To do this, the nearest noble gas that precedes the element in question is written first, and then the electron configuration is 2

continued from that point forward. For example, the electron notation of carbon is 1s 2s 2p , 2

and the noble gas notation is [He]2s 2p . This notation makes it easier to identify elements, and is shorter than writing out the full notation of atomic orbitals. Compounds

Structure of XeF4, one of the first noble gas compounds to be discovered

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The noble gases show extremely low chemical reactivity; consequently, only a few hundred noble gas compounds have been formed. Neutral compounds in which helium and neon are involved in chemical bonds have not been formed (although there is some theoretical evidence for a few helium compounds), while xenon, krypton, and argon have shown only minor reactivity. The reactivity follows the order Ne < He < Ar < Kr < Xe < Rn. In 1933, Linus Pauling predicted that the heavier noble gases could form compounds with fluorine and oxygen. He predicted the existence of krypton hexafluoride (KrF 6) and xenon hexafluoride (XeF6), speculated that XeF8 might exist as an unstable compound, and suggested xenic acid could form perxenate salts. These predictions were shown to be generally accurate, except that XeF8 is now thought to be both thermodynamically and kinetically unstable. Xenon compounds are the most numerous of the noble gas compounds that have been formed. Most of them have the xenon atom in the oxidation state of +2, +4, +6, or +8 bonded to highly electronegative atoms such as fluorine or oxygen, as in xenon difluoride (XeF2), xenon tetrafluoride (XeF4), xenon hexafluoride (XeF6), xenon tetroxide (XeO4), and sodium perxenate (Na4XeO6). Some of these compounds have found use in chemical synthesis as oxidizing agents; XeF2, in particular, is commercially available and can be used as a fluorinating agent. As of 2007, about five hundred compounds of xenon bonded to other elements have been identified, including organoxenon compounds (containing xenon bonded to carbon), and xenon bonded to nitrogen, chlorine, gold, mercury, and xenon itself. Compounds of xenon bound to boron, hydrogen, bromine, iodine, beryllium, sulphur, titanium, copper, and silver have also been observed but only at low temperatures in noble gas matrices, or in supersonic noble gas jets. In theory, radon is more reactive than xenon, and therefore should form chemical bonds more easily than xenon does. However, due to the high radioactivity and short half-life of radon isotopes, only a few fluorides and oxides of radon have been formed in practice. Krypton is less reactive than xenon, but several compounds have been reported with krypton in the oxidation state of +2.Krypton difluoride is the most notable and easily characterized. Compounds in which krypton forms a single bond to nitrogen and oxygen have also been characterized, but are only stable below −60 °C (−76 °F) and −90 °C (−130 °F) respectively. Krypton atoms chemically bound to other nonmetals (hydrogen, chlorine, carbon) as well as some late transition metals (copper, silver, gold) have also been observed, but only either at low temperatures in noble gas matrices, or in supersonic noble gas jets. Similar conditions were used to obtain the first few compounds of argon in 2000, such as argon fluorohydride (HArF), and some bound to the late transition metals copper, silver, and gold. As of 2007, no

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stable neutral molecules involving covalently bound helium or neon are known. The noble gasesâ&#x20AC;&#x201D;including heliumâ&#x20AC;&#x201D;can form stable molecular ions in the gas phase. The +

simplest is the helium hydride molecular ion, HeH , discovered in 1925. Because it is composed of the two most abundant elements in the universe, hydrogen and helium, it is believed to occur naturally in the interstellar medium, although it has not been detected yet. In addition to these ions, there are many known neutral excimers of the noble gases. These are compounds such as ArF and KrF that are stable only when in an excited electronic state; some of them find application in excimer lasers. In addition to the compounds where a noble gas atom is involved in a covalent bond, noble gases also form non-covalent compounds. The clathrates, first described in 1949, consist of a noble gas atom trapped within cavities of crystal lattices of certain organic and inorganic substances. The essential condition for their formation is that the guest (noble gas) atoms must be of appropriate size to fit in the cavities of the host crystal lattice. For instance, argon, krypton, and xenon form clathrates with hydroquinone, but helium and neon do not because they are too small or insufficiently polarizable to be retained. Neon, argon, krypton, and xenon also form clathrate hydrates, where the noble gas is trapped in ice.

An endohedral fullerene compound containing a noble gas atom Noble gases can form endohedral fullerene compounds, in which the noble gas atom is trapped inside a fullerene molecule. In 1993, it was discovered that when C 60, a spherical molecule consisting of 60 carbon atoms, is exposed to noble gases at high pressure, 60

can

complexes be

formed

(the

such notation

indicates

as He

contained

He@C inside

60 but not covalently bound to it). As of 2008, endohedral complexes with helium, neon, argon, krypton, and xenon have been obtained. These compounds have found use in the study of the structure and reactivity of fullerenes by means of the nuclear magnetic

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resonance of the noble gas atom.

Bonding in XeF2 according to the 3-center-4-electron bond model Noble gas compounds such as xenon difluoride (XeF 2) are considered to be hypervalent because they violate the octet rule. Bonding in such compounds can be explained using a three-center four-electron bond model. This model, first proposed in 1951, considers bonding of three collinear atoms. For example, bonding in XeF 2 is described by a set of three molecular orbitals (MOs) derived from p-orbitals on each atom. Bonding results from the combination of a filled p-orbital from Xe with one half-filled p-orbital from each F atom, resulting in a filled bonding orbital, a filled non-bonding orbital, and an empty antibonding orbital. The highest occupied molecular orbital is localized on the two terminal atoms. This represents a localization of charge which is facilitated by the high electronegativity of fluorine. The chemistry of heavier noble gases, krypton and xenon, are well established. The chemistry of the lighter ones, argon and helium, is still at an early stage, while a neon compound is still yet to be identified. Occurrence and production The abundances of the noble gases in the universe decrease as their atomic numbers increase. Helium is the most common element in the universe after hydrogen, with a mass fraction of about 24%. Most of the helium in the universe was formed during Big Bang nucleosynthesis, but the amount of helium is steadily increasing due to the fusion of hydrogen in stellar nucleosynthesis (and, to a very slight degree, the alpha decay of heavy elements). Abundances on Earth follow different trends; for example, helium is only the third most abundant noble gas in the atmosphere. The reason is that there is no primordial helium in the atmosphere; due to the small mass of the atom, helium cannot be retained by the Earth's gravitational field. Helium on Earth comes from the alpha decay of heavy elements such as uranium and thorium found in the Earth's crust, and tends to accumulate in natural gas deposits. The abundance of argon, on the other hand, is increased as a result of the beta decay of potassium-40, also found in the Earth's crust, to form argon-40, which is the

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most abundant isotope of argon on Earth despite being relatively rare in the Solar System. This process is the base for the potassium-argon dating method. Xenon has an unexpectedly low abundance in the atmosphere, in what has been called the missing xenon problem; one theory is that the missing xenon may be trapped in minerals inside the Earth's crust. After the discovery of xenon dioxide, a research showed that Xe can substitute for Si in the quartz. Radon is formed in the lithosphere as from the alpha decay of radium. It can seep into buildings through cracks in their foundation and accumulate in areas that are not well ventilated. Due to its high radioactivity, radon presents a significant health hazard; it is implicated in an estimated 21,000 lung cancer deaths per year in the United States alone. Abundance

Helium Neon

Solar System (for each atom of silicon) 2343

Earth's atmosphere (volume fraction in ppm) Igneous rock (mass fraction in ppm)

5.20 −3

Argon Krypton 5.515

2.148

0.1025

18.20

9340.00 1.10 −5

Xenon ×5.391

−5

−6

0.09

−2

3 × 10 7 × 10 4 × 10 –

–

Radon ×

–

(0.06–18) −19

1.7 × 10

Gas

2004 price (USD/m )

Helium (industrial grade)

4.20–4.90

Helium (laboratory grade)

22.30–44.90

Argon

2.70–8.50

Neon

60–120

Krypton

400–500

Xenon

4000–5000

Neon, argon, krypton, and xenon are obtained from air using the methods of liquefaction of gases, to convert elements to a liquid state, and fractional distillation, to separate mixtures into component parts. Helium is typically produced by separating it from natural gas, and radon is isolated from the radioactive decay of radium compounds. The prices of the noble gases are influenced by their natural abundance, with argon being the cheapest and xenon the most expensive. As an example, the table to the right lists the 2004 prices in the United States for laboratory quantities of each gas.

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−10

Applications

Liquid helium is used to cool the superconducting magnets in modern MRI scanners. Noble gases have very low boiling and melting points, which makes them useful as cryogenic refrigerants. In particular, liquid helium, which boils at 4.2 K (−268.95 °C; −452.11 °F), is used for superconducting magnets, such as those needed in nuclear magnetic resonance imaging and nuclear magnetic resonance. Liquid neon, although it does not reach temperatures as low as liquid helium, also finds use in cryogenics because it has over 40 times more refrigerating capacity than liquid helium and over three times more than liquid hydrogen. Helium is used as a component of breathing gases to replace nitrogen, due its low solubility in fluids, especially in lipids. Gases are absorbed by the blood and body tissues when under pressure like in scuba diving, which causes an anesthetic effect known as nitrogen narcosis. Due to its reduced solubility, little helium is taken into cell membranes, and when helium is used to replace part of the breathing mixtures, such as in trimix or heliox, a decrease in the narcotic effect of the gas at depth is obtained. Helium's reduced solubility offers further advantages for the condition known as decompression sickness, or the bends. The reduced amount of dissolved gas in the body means that fewer gas bubbles form during the decrease in pressure of the ascent. Another noble gas, argon, is considered the best option for use as a drysuit inflation gas for scuba diving. Helium is also used as filling gas in nuclear fuel rods for nuclear reactors.

Goodyear Blimp Since the Hindenburg disaster in 1937, helium has replaced hydrogen as a lifting gas in blimps and balloons due to its lightness and incombustibility, despite an 8.6% decrease in

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buoyancy. In many applications, the noble gases are used to provide an inert atmosphere. Argon is used in the synthesis of air-sensitive compounds that are sensitive to nitrogen. Solid argon is also used for the study of very unstable compounds, such as reactive intermediates, by trapping them in an inert matrix at very low temperatures. Helium is used as the carrier medium in gas chromatography, as a filler gas for thermometers, and in devices for measuring radiation, such as the Geiger counter and the bubble chamber. Helium and argon are both commonly used to shield welding arcs and the surrounding base metal from the atmosphere during welding and cutting, as well as in other metallurgical processes and in the production of silicon for the semiconductor industry.

15,000-watt xenon short-arc lamp used in IMAX projectors Noble gases are commonly used in lighting because of their lack of chemical reactivity. Argon, mixed with nitrogen, is used as a filler gas for incandescent light bulbs. Krypton is used in high-performance light bulbs, which have higher color temperatures and greater efficiency, because it reduces the rate of evaporation of the filament more than argon; halogen lamps, in particular, use krypton mixed with small amounts of compounds of iodine or bromine. The noble gases glow in distinctive colors when used inside gas-discharge lamps, such as "neon lights". These lights are called after neon but often contain other gases and phosphors, which add various hues to the orange-red color of neon. Xenon is commonly used in xenon arc lamps which, due to their nearly continuous spectrum that resembles daylight, find application in film projectors and as automobile headlamps. The noble gases are used in excimer lasers, which are based on short-lived electronically excited molecules known as excimers. The excimers used for lasers may be noble gas dimers such as Ar2, Kr2 or Xe2, or more commonly, the noble gas is combined with a halogen in excimers such as ArF, KrF, XeF, or XeCl. These lasers produce ultraviolet light which, due to its short wavelength (193 nm for ArF and 248 nm for KrF), allows for high-precision imaging. Excimer lasers have many industrial, medical, and scientific applications. They are used for microlithography and microfabrication, which are essential for integrated circuit

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manufacture, and for laser surgery, including laser angioplasty and eye surgery. Some noble gases have direct application in medicine. Helium is sometimes used to improve the ease of breathing of asthma sufferers. Xenon is used as an anesthetic because of its high solubility in lipids, which makes it more potent than the usual nitrous oxide, and because it is readily eliminated from the body, resulting in faster recovery. Xenon finds application in medical imaging of the lungs through hyperpolarized MRI. Radon, which is highly radioactive and is only available in minute amounts, is used in radiotherapy. Discharge color Colors and spectra (bottom row) of electric discharge in noble gases; only the second row represents pure gases.

Helium

Neon

Argon

Krypton

Xenon

(with some Hg in the "Ar" image)

The color of gas discharge emission depends on several factors, including the following: 

discharge parameters (local value of current density and electric field, temperature, etc. – note the color variation along the discharge in the top row);



gas purity (even small fraction of certain gases can affect color);



material of the discharge tube envelope – note suppression of the UV and blue

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components in the bottom-row tubes made of thick household glass. See also 

Noble gas (data page), for extended tables of physical properties.



Noble metal, for metals that are resistant to corrosion or oxidation.



Inert gas, for any gas that is not reactive under normal circumstances.



Industrial gas



Neutronium



Noble gas configuration

References 

Bennett, Peter B.; Elliott, David H. (1998). The Physiology and Medicine of Diving. SPCK Publishing. ISBN 0-7020-2410-4.



Bobrow Test Preparation Services (2007-12-05). CliffsAP Chemistry. CliffsNotes. ISBN 0-470-13500-X.



Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.



Harding, Charlie J.; Janes, Rob (2002). Elements of the P Block. Royal Society of Chemistry. ISBN 0-85404-690-9.



Holloway, John H. (1968). Noble-Gas Chemistry. London: Methuen Publishing. ISBN 0-412-21100-9.



Mendeleev, D. (1902–1903). Osnovy Khimii (The Principles of Chemistry) (in Russian) (7th ed.).



{ {cite book|first=Minoru|last=Ojima|coauthors=Podosek, Frank A.|title=Noble Gas Geochemistry|year=2002|publisher=Cambridge University Press|isbn=0-521-803667|url=http://books.google.com/?id=CBM2LJDvRtgC%7Cref=CITEREFOjima2002}}



Weinhold, F.; Landis, C. (2005). Valency and bonding. Cambridge University Press. ISBN 0-521-83128-8.



Scerri, Eric R. (2007). The Periodic Table, Its Story and Its Significance. Oxford University Press. ISBN 0-19-530573-6.

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Article Sources and Contributors 

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