Chemistry (Take to Phone)

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Fajans' rules

In inorganic chemistry,Fajans' Rules, formulated by Kazimierz Fajans in 1923,[1][2][3]are used to predict

whether a chemical bond will be covalent orionic, and depend on the charge on the cation and the

relative sizes of the cation and anion. They can be summarized in the following table:

Chart illustrating the relationship between atomic and ionic radius

Ionic Covalent

Low positive charge High positive charge

Large cation Small cation

Small anion Large anion

Thus sodium chloride (with a low positive charge (+1), a fairly large cation (~1 ) and relatively small

anion (2 ) is ionic; but aluminium iodide (AlI3) (with a high positive charge (+3) and a large anion) is

covalent.

Polarization will be increased by:

High charge and small size of the cation

Ionic potential Z+/r+ (= polarizing power)

High charge and large size of the anion

The polarizability of an anion is related to the deformability of its electron cloud (i.e. its

"softness")

An incomplete valence shell electron configuration

Noble gas configuration of the cation produces better shielding and less polarizing power

e.g. Hg2+ (r+ = 102 pm) is more polarizing than Ca2+ (r+ = 100 pm)

Considering an ionic bond, in this the electron(s) is(are) not shared but transferred between the

atoms conveying definite charges to each participant in the bond. The "size" of charge of the charge

depends on the number of electrons transferred so an aluminium atom with a +3 charge has a

relatively large positive charge. That positive charge then exerts an attractive force on the electron

cloud of the other ion, which has accepted the electrons from the aluminium (or other) positive ion.

Two contrasting examples can illustrate the variation in effects. In the case of aluminium iodide an

ionic bond with much covalent character is present. In the AlI3 bonding, the aluminium gains a +3

charge. The large charge pulls on the electron cloud of the iodines. Now, if we consider the iodine

atom, we see that it is relatively large and thus the outer shell electrons are relatively well shielded
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from the nuclear charge. In this case, the aluminium ion's charge will "tug" on the electron cloud of

iodine, drawing it closer to itself. As the electron cloud of the iodine nears the aluminium atom, the

negative charge of the electron cloud "cancels" out the positive charge of the aluminium cation. This

produces an ionic bond with covalent character.

Now, if we take a different example, for example AlF3. We see that a similar situation occurs, but

instead of iodine we now have fluorine, a relatively small highly electronegative atom. The fluorine's

electron cloud is definitely less shielded from the nuclear charge and will thus be less polarizable.

Thus, we get an ionic compound (metal bonded to a nonmetal) with slight covalent character.

Boron group

Characteristics[edit]

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 pershell

5 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.[5]
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Chemical reactivity[edit]

Hydrides[edit]

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 calledboranes.[6]The simplest borane is diborane, or B2H6.[5]Another example

is B10H14.

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 H3InP(Cy)3, which is considered to be a phosphine.[7]No stable compound of

thallium and hydrogen has been synthesized in any laboratory.

[show]Some common chemical compounds of the boron group[5]

[8][9][10][11][12]

Oxides[edit]

All of the boron-group elements are known to form a trivalent oxide, with two atoms of the element

bonded covalently with three atoms ofoxygen. These elements show a trend of

increasing pH (from acidic to basic).[13] Boron oxide (B2O3) is slightly acidic, aluminium andgallium

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.

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

Halides[edit]

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.[14] Fluorine, the first halogen, is able to form stable compounds with every element

that has been tested (except neon and helium),[15]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
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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 2AlI3.[16] Astatine, the heaviest halogen, has

only formed a few compounds, due to its radioactivity and short half-life, and no reports of a compound

with an AtB,Al,Ga,In,Tl, orUut bond have been seen, although scientists think that it shouldform salts with metals.[17]

Physical properties[edit]

It has been noticed that the elements in the boron group have similarphysical 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 theirvalence 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.[8]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 betterthan most non-metals.[18]

Oxidation states[edit]

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.[19]Boron is capable of forming compounds with lower oxidization states, of +1 or +2, and

aluminium can do the same.[20]Gallium cannot form compounds with the oxidation state +2 but can with

+1 and +3. Indium is like gallium, but its +1 compounds are more stable than those of the other 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[edit]

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.

Element Boiling Point (C) Density (g/cm3)

Boron 4,000 2.46

Aluminium 2,519 2.7

Gallium 2,204 5.904
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Indium 2,072 7.31

Thallium 1,473 11.85

Nuclear[edit]

With the exception of the synthetic ununtrium, all of the elements of the boron group have stable isotopes.

Because all theiratomic numbers are odd, boron, gallium and thallium have only two stable isotopes,

while aluminium and indium are monoisotopic, having only one. 10B and 11B are both stable, as

are 27Al, 69Ga and 71Ga, 113In, and 203Tl and205Tl.[21]All of these isotopes are readily found in macroscopic

quantities in nature. In theory, though, all isotopes with an atomic numbergreater 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 ofproton decay, which has never been observed).

Like all other elements, the elements of the boron group have radioactive isotopes, either foundin trace quantities in nature or produced synthetically. The longest-lived of these unstable isotopes is

the indium isotope 115In, with its extremely long half-life of 4.41 1014 y. This isotope is relatively

important among indium's radioisotopes. The shortest-lived is7B, with a half-life of a mere 35050

1024 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.

Carbon group


Chemical[edit]



Z Element No. of electrons/shell

6 Carbon 2, 4

14 Silicon 2, 8, 4

32 Germanium 2, 8, 18, 4
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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 outerenergy level. The last orbital of all

these elements is the p2 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 increasing atomic

number. Carbon alone forms negative ions, in the form ofcarbide (C4) 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.[1]

Silicon forms two hydrides: SiH4and Si2H6. Silicon forms tetrahalides with fluorine, chlorine, and iodine.

Silicon also forms a dioxide anda disulfide.[2] Silicon nitride has the formula Si3N4.[3]

Germanium forms two hydrides: GeH4and 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 Ge3N4.[4]

Tin forms two hydrides: SnH4and Sn2H6. Tin forms dihalides and tetrahalides with all halogens exceptastatine. 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.[5]

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.[6]

There are no known compounds of flerovium.[7]

Physical[edit]

The boiling points of the carbon group tend to get lower with the heavier elements. Carbon, the lightest

carbon group element, sublimatesat 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.[8]

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.[8]
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kipedia.org/wiki/Carbon_group#cite_note-5https://en.wikipedia.org/w/index.php?title=Distannane&action=edit&redlink=1https://en.wikipedia.org/wiki/Stannanehttps://en.wikipedia.org/wiki/Carbon_group#cite_note-4https://en.wikipedia.org/wiki/Germanium_nitridehttps://en.wikipedia.org/wiki/Germanium_nitridehttps://en.wikipedia.org/wiki/Digermanehttps://en.wikipedia.org/wiki/Germanehttps://en.wikipedia.org/wiki/Carbon_group#cite_note-The_Elements-3https://en.wikipedia.org/wiki/Silicon_nitridehttps://en.wikipedia.org/wiki/Carbon_group#cite_note-2https://en.wikipedia.org/wiki/Silicon_disulfidehttps://en.wikipedia.org/wiki/Silicon_dioxidehttps://en.wikipedia.org/wiki/Disilanehttps://en.wikipedia.org/wiki/Silanehttps://en.wikipedia.org/wiki/Carbon_group#cite_note-1https://en.wikipedia.org/wiki/Carbon_dioxidehttps://en.wikipedia.org/wiki/Carbon_suboxidehttps://en.wikipedia.org/wiki/Carbon_monoxidehttps://en.wikipedia.org/wiki/Carbon_monoxidehttps://en.wikipedia.org/wiki/Astatinehttps://en.wikipedia.org/wiki/Halogenshttps://en.wikipedia.org/wiki/Noble_gashttps://en.wikipedia.org/wiki/Radioactivehttps://en.wikipedia.org/wiki/Metalhttps://en.wikipedia.org/wiki/Leadhttps://en.wikipedia.org/wiki/Tinhttps://en.wikipedia.org/wiki/Metalloidshttps://en.wikipedia.org/wiki/Germaniumhttps://en.wikipedia.org/wiki/Siliconhttps://en.wikipedia.org/wiki/Carbidehttps://en.wikipedia.org/wiki/Ionhttps://en.wikipedia.org/wiki/Carbonhttps://en.wikipedia.org/wiki/Atomhttps://en.wikipedia.org/wiki/Atomic_orbitalhttps://en.wikipedia.org/wiki/Chemical_elementhttps://en.wikipedia.org/wiki/Energy_levelhttps://en.wikipedia.org/wiki/Electronhttps://en.wikipedia.org/wiki/Fleroviumhttps://en.wikipedia.org/wiki/Leadhttps://en.wikipedia.org/wiki/Tin


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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.[8]

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.[8]

Al lo tropes

Carbon has mu l t ip leal lotropes. The most common isgraphite, which is carbon in the form of

stacked sheets. Another form o f carbon isd iamond , but this is relatively

rare.Amorphouscarbonis a th i rd al lo trope of carbon; i t is a com ponent ofsoo t . Another al lo trope

of carbon is aful lerene, which h as the form of s heets of carbon atoms fo lded into a sphere. A f i f th

al lotrope of c arbon, dis cov ered in 2003, is cal ledgraphene, and is in the form of a layer of carbon

atoms arranged in a honeycomb-shaped format ion.[3][9][ 10]

Silicon has two known allotropes that exist at room temperature. These allotropes are known as theamorphous and the crystalline allotropes. The amorphous allotrope is a brown powder. The crystalline

allotrope is gray and has a metallic luster.[11]

Tin has two allotropes: -tin, also known as gray tin, and -tin. Tin is typically found in the -tin form, a

silvery metal. However, at standard pressure, -tin converts to -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.[3][12]

Nuclear[edit]

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.[12]

Isotopes[edit]

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.[8] Carbon-14 is a natural radioactive isotope with a half-

life of 5,730 years.[13]

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 ofactinides, and via spallation in the upper

atmosphere. Silicon-34 also occurs naturally as the result of radioactive decay of actinides.[13]

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.[13]

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
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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.[13]

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.[13]

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.

Pnictogen


Chemical[edit]




7 nitrogen 2, 5

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 arenitrogen(N), which in its diatomic form is the principal component of

air, andphosphorus(P), which, like nitrogen, is essential to all known forms of life.
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Compounds[edit]

Binary compounds of the group can be referred to collectively as pnictides. The spelling derives from

the Greek verb (pngein), 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 ,pnte, five) was also used for this group at one time,[4]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 pnictides. These are in the

form of REaMbPnc, where M is a carbon group orboron group element and Pn is any pnictogen except

nitrogen. These compounds are between ionic and covalent compounds and thus have unusual bonding

properties.[5]

These elements are also noted for theirstability 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 anothernoble gas would be too expensive.

The upper pnictogens, that is, nitrogen, phosphorus, and arsenic tend to form -3 charges. Antimony can

either take on a +3 or +5, by losing its p-shell electrons or losing its p-shell and s-shell electrons,

respectively.

Chalcogen

The chalcogens(/klkdnz/) are the chemical elements in group 16 of the periodic table. This group

is also known as the oxygen family orgroup 16. It consists of the

elements oxygen (O), sulfur(S), selenium (Se), tellurium (Te), and the radioactive elementpolonium (Po).

The synthetic element livermorium (Lv) is predicted to be a chalcogen as well.[1]The word chalcogen

comes from the Greek word chalkos, meaning "bronze" or "ore", and the word gens, meaning "born".[2][3]

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.[4]

Lighter chalcogens are typically nontoxic in their elemental form, and are often critical to life, while the

heavier chalcogens are typicallytoxic.[1]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.[5]Tellurium often
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ia.org/wiki/Mnemonichttps://en.wikipedia.org/wiki/Chokinghttps://en.wikipedia.org/wiki/Greek_languagehttps://en.wikipedia.org/w/index.php?title=Pnictogen&action=edit&section=3


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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 chemical industry.[5]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.[1]

Properties[edit]

Atomic and physical[edit]

Chalcogens show similar patterns in electron configuration, especially in the outermost shells, where they

all have the same number ofvalence electrons, resulting in similar trends in chemical behavior:


8 oxygen 2, 6

16 sulfur 2, 8, 6

34 selenium 2, 8, 18, 6

52 tellurium 2, 8, 18, 18, 6

84 polonium 2, 8, 18, 32, 18, 6

116 livermorium 2, 8, 18, 32, 32, 18, 6 (predicted)[6]
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Element Melting point (Celsius) Boiling point (Celsius) Reference

Oxygen -219 -183 [4]

Sulfur 120 445 [4]

Selenium 221 685 [4]

Tellurium 450 988 [4]

Pollonium 254 962 [4]

Element Grams per cubic centimeter at STP Reference

Oxygen 0.00143 [4]

Sulfur 2.07 [4]

Selenium 4.3 [4]

Tellurium 6.24 [4]

Polonium 9.2 [4]

All chalcogens have six valence electrons. Most of the solid chalcogens are soft[7]and do not conductheat well.[4] Electronegativitydecreases towards the chalcogens with higher atomic numbers. Density,

melting and boiling points, and atomic and ionic radii[8]tend to increase towards the chalcogens with

higher atomic numbers.[4]

Isotopes[edit]

Out of the six known chalcogens, one (oxygen) has an atomic number equal to a nuclearmagic number,

which means that theiratomic nuclei tend to have increased stability towards radioactive decay.[9]Oxygen
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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.[10]It has an additional 28 isomers.[1]In addition to the

stable isotopes, some radioactive chalcogen isotopes occur in nature, either because they are decay

products, such as210

Po, because they are primordial, such as82

Se, because ofcosmic ray spallation, orvia nuclear fission of uranium. Livermorium isotopes 290 through 293 have been discovered. The most

stable livermorium isotope is 293Lv, which has a half-life of 0.061 seconds.[1][11]

Among the lighter chalcogens (oxygen and sulfur), the most neutron-starved isotopes undergo proton

emission, the moderately neutron-starved isotopes undergo electron capture or+ decay, the moderately

neutron-rich isotopes undergo - 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.[12]Isotopes

with nuclear spins are more common among the chalcogens selenium and tellurium than they are with

sulfur.[

Halogen


Chemical[edit]

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 X2 HX BX3 AlX3 CX4

F 159 574 645 582 456

Cl 243 428 444 427 327

Br 193 363 368 360 272

I 151 294 272 285 239

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 atoms being highly electronegative 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 heaviernoble 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
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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.

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[edit]

Diatomic halogen molecules[edit]

The halogens form homonucleardiatomic molecules (not proven for astatine). As such they form part of

the group known as "elemental gases".

halogen molecule structure modeld(XX) / pm

(gas phase)

d(XX) / pm

(solid phase)

fluorine F2 143 149

chlorine Cl2 199 198

bromine Br2 228 227

iodine I2 266 272

astatine At2

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

Compounds[edit]

Hydrogen halides[edit]

All of the halogens except for astatine have been observed to react with hydrogen to form hydrogen

halides. For fluorine, chlorine, and bromine, this reaction is in the form of:
http://en.wikipedia.org/wiki/Silicon_tetrafluoridehttp://en.wikipedia.org/wiki/Polytetrafluoroethylenehttp://en.wikipedia.org/wiki/Organofluorinehttp://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=5http://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=6http://en.wikipedia.org/wiki/Homonuclearhttp://en.wikipedia.org/wiki/Diatomichttp://en.wikipedia.org/wiki/Moleculeshttp://en.wikipedia.org/wiki/Fluorinehttp://en.wikipedia.org/wiki/Fluorinehttp://en.wikipedia.org/wiki/Chlorinehttp://en.wikipedia.org/wiki/Chlorinehttp://en.wikipedia.org/wiki/Brominehttp://en.wikipedia.org/wiki/Iodinehttp://en.wikipedia.org/wiki/Iodinehttp://en.wikipedia.org/wiki/Astatinehttp://en.wikipedia.org/wiki/Astatinehttp://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=7http://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=8http://en.wikipedia.org/wiki/File:Iodine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Diiodine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Bromine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dibromine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Chlorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dichlorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Fluorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Difluorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Iodine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Diiodine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Bromine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dibromine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Chlorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dichlorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Fluorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Difluorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Iodine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Diiodine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Bromine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dibromine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Chlorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dichlorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Fluorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Difluorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Iodine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Diiodine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Bromine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dibromine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Chlorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dichlorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Fluorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Difluorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Iodine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Diiodine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Bromine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dibromine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Chlorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dichlorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Fluorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Difluorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Iodine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Diiodine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Bromine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dibromine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Chlorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dichlorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Fluorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Difluorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Iodine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Diiodine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Bromine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dibromine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Chlorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dichlorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Fluorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Difluorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Iodine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Diiodine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Bromine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dibromine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Chlorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Dichlorine-2D-dimensions.pnghttp://en.wikipedia.org/wiki/File:Fluorine-3D-vdW.pnghttp://en.wikipedia.org/wiki/File:Difluorine-2D-dimensions.pnghttp://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=8http://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=7http://en.wikipedia.org/wiki/Astatinehttp://en.wikipedia.org/wiki/Iodinehttp://en.wikipedia.org/wiki/Brominehttp://en.wikipedia.org/wiki/Chlorinehttp://en.wikipedia.org/wiki/Fluorinehttp://en.wikipedia.org/wiki/Moleculeshttp://en.wikipedia.org/wiki/Diatomichttp://en.wikipedia.org/wiki/Homonuclearhttp://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=6http://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=5http://en.wikipedia.org/wiki/Organofluorinehttp://en.wikipedia.org/wiki/Polytetrafluoroethylenehttp://en.wikipedia.org/wiki/Silicon_tetrafluoride


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H2 + X2 2HX

However, hydrogen iodide can split back into its constituent elements.[5]

The hydrogen-halogen reactions get gradually less reactive towards 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 only explosive when exposed to flames. Iodine only partially reacts with hydrogen, forming

an equilibrium.[5]

The halogens all 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, and hydroiodic acid are

all strong acids, but hydrofluoric acid is a weak acid.[6]

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, causingpulmonary edema and

damaging cells.[7]Hydrogen chloride is also a dangerous chemical. Breathing in gas with more than fiftyparts per million of hydrogen chloride can cause death in humans.[8]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.[9]Hydrogen iodide, like other hydrogen halides, is toxic.[10]

Metal halides[edit]

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[5]

Iron reacts with fluorine, chlorine, and bromine to form Iron(III) halides. These reactions are in the form of:

2Fe + 3X2 2FeX3[5]

However, when iron reacts with iodine, it only forms iron(II) iodide.

Iron wool can react rapidly with f luorine 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 will reaction will instead result in a reddish-

brown product. Iron can also react with bromine to formiron(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 thereaction 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.[5]

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
http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Chemical_equilibriumhttp://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Hydrogen_halidehttp://en.wikipedia.org/wiki/Hydrogen_fluoridehttp://en.wikipedia.org/wiki/Hydrogen_fluoridehttp://en.wikipedia.org/wiki/Hydrogen_chloridehttp://en.wikipedia.org/wiki/Hydrogen_bromidehttp://en.wikipedia.org/wiki/Hydrogen_iodidehttp://en.wikipedia.org/wiki/Hydrogen_astatidehttp://en.wikipedia.org/wiki/Hydrogen_astatidehttp://en.wikipedia.org/wiki/Hydrogen_bondhttp://en.wikipedia.org/wiki/Strong_acidhttp://en.wikipedia.org/wiki/Weak_acidhttp://en.wikipedia.org/wiki/Halogen#cite_note-6http://en.wikipedia.org/wiki/Halogen#cite_note-6http://en.wikipedia.org/wiki/Halogen#cite_note-6http://en.wikipedia.org/wiki/Irritationhttp://en.wikipedia.org/wiki/Acidhttp://en.wikipedia.org/wiki/Industryhttp://en.wikipedia.org/wiki/Pulmonary_edemahttp://en.wikipedia.org/wiki/Halogen#cite_note-7http://en.wikipedia.org/wiki/Halogen#cite_note-7http://en.wikipedia.org/wiki/Halogen#cite_note-7http://en.wikipedia.org/wiki/Halogen#cite_note-8http://en.wikipedia.org/wiki/Halogen#cite_note-8http://en.wikipedia.org/wiki/Halogen#cite_note-8http://en.wikipedia.org/wiki/Halogen#cite_note-9http://en.wikipedia.org/wiki/Halogen#cite_note-9http://en.wikipedia.org/wiki/Halogen#cite_note-9http://en.wikipedia.org/wiki/Halogen#cite_note-10http://en.wikipedia.org/wiki/Halogen#cite_note-10http://en.wikipedia.org/wiki/Halogen#cite_note-10http://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=9http://en.wikipedia.org/wiki/Metal_halideshttp://en.wikipedia.org/wiki/Sodium_fluoridehttp://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Sodium_bromidehttp://en.wikipedia.org/wiki/Sodium_bromidehttp://en.wikipedia.org/wiki/Sodium_iodidehttp://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Iron(II)_iodidehttp://en.wikipedia.org/wiki/Iron(III)_fluoridehttp://en.wikipedia.org/wiki/Iron_(III)_chloridehttp://en.wikipedia.org/wiki/Iron_(III)_chloridehttp://en.wikipedia.org/wiki/Iron(III)_bromidehttp://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Iron(III)_bromidehttp://en.wikipedia.org/wiki/Iron_(III)_chloridehttp://en.wikipedia.org/wiki/Iron_(III)_chloridehttp://en.wikipedia.org/wiki/Iron(III)_fluoridehttp://en.wikipedia.org/wiki/Iron(II)_iodidehttp://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Sodium_iodidehttp://en.wikipedia.org/wiki/Sodium_bromidehttp://en.wikipedia.org/wiki/Sodium_bromidehttp://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Sodium_fluoridehttp://en.wikipedia.org/wiki/Metal_halideshttp://en.wikipedia.org/w/index.php?title=Halogen&action=edit&section=9http://en.wikipedia.org/wiki/Halogen#cite_note-10http://en.wikipedia.org/wiki/Halogen#cite_note-9http://en.wikipedia.org/wiki/Halogen#cite_note-8http://en.wikipedia.org/wiki/Halogen#cite_note-7http://en.wikipedia.org/wiki/Pulmonary_edemahttp://en.wikipedia.org/wiki/Industryhttp://en.wikipedia.org/wiki/Acidhttp://en.wikipedia.org/wiki/Irritationhttp://en.wikipedia.org/wiki/Halogen#cite_note-6http://en.wikipedia.org/wiki/Weak_acidhttp://en.wikipedia.org/wiki/Strong_acidhttp://en.wikipedia.org/wiki/Hydrogen_bondhttp://en.wikipedia.org/wiki/Hydrogen_astatidehttp://en.wikipedia.org/wiki/Hydrogen_astatidehttp://en.wikipedia.org/wiki/Hydrogen_iodidehttp://en.wikipedia.org/wiki/Hydrogen_bromidehttp://en.wikipedia.org/wiki/Hydrogen_chloridehttp://en.wikipedia.org/wiki/Hydrogen_fluoridehttp://en.wikipedia.org/wiki/Hydrogen_fluoridehttp://en.wikipedia.org/wiki/Hydrogen_halidehttp://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5http://en.wikipedia.org/wiki/Chemical_equilibriumhttp://en.wikipedia.org/wiki/Halogen#cite_note-assorted-5


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ClF3 can be produced by a reaction of a pure halogen with a smaller interhalogen such as ClF. All

interhalogens except IF7can be produced by directly combining pure halogens in various conditions.[11]

Interhalogens are typically more reactive than all diatomic halogen molecules except F2 because

interhalogen bonds are weaker. However, the chemical properties of interhalogens are still roughly the

same as those ofdiatomic 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 BrF3, and many

iodine-containing interhalogens are solids.[11]

Organohalogen compounds

Many synthetic organic compounds such as plastic polymers, and a few natural ones, contain halogen

atoms; these are known as halogenatedcompounds ororganic 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 transmitterGABA and are also used by the body to produce stomach acid. Iodine is needed in

trace amounts for the production ofthyroid hormones such as thyroxine. On the other hand, neither

fluorine nor bromine are believed to be essential for humans. Organohalogens are also synthesized

through the nucleophilic abstraction reaction.
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