Iodine

[Guest guest]

Iodine

From Wikipedia, the free encyclopedia

Jump to: navigation, search

53 tellurium & #8592; iodine & #8594; xenon

Br

& #8593;

I

& #8595;

At

Periodic Table - Extended Periodic Table

General

Name, Symbol, Number iodine, I, 53

Chemical series halogens

Group, Period, Block 17, 5, p

Appearance violet-dark gray, lustrous

Standard atomic weight 126.90447(3)

& #8201;g·mol & #8722;1

Electron configuration [Kr] 4d10 5s2 5p5

Electrons per shell 2, 8, 18, 18, 7

Physical properties

Phase solid

Density (near r.t.) 4.933 & #8201;g·cm & #8722;3

Melting point 386.85 & #8201;K

(113.7 & #8201;°C, 236.66 & #8201;°F)

Boiling point 457.4 & #8201;K

(184.3 & #8201;°C, 363.7 & #8201;°F)

Critical point 819 K, 11.7 MPa

Heat of fusion (I2) 15.52 & #8201;kJ·mol & #8722;1

Heat of vaporization (I2) 41.57 & #8201;kJ·mol & #8722;1

Heat capacity (25 & #8201;°C) (I2) 54.44

& #8201;J·mol & #8722;1·K & #8722;1

Vapor pressure (rhombic) P(Pa) 1 10 100 1 k 10 k 100 k

at T(K) 260 282 309 342 381 457

Atomic properties

Crystal structure orthorhombic

Oxidation states ±1, 5, 7

(strongly acidic oxide)

Electronegativity 2.66 (ing scale)

Ionization energies 1st: 1008.4 kJ/mol

2nd: 1845.9 kJ/mol

3rd: 3180 kJ/mol

Atomic radius 140 & #8201;pm

Atomic radius (calc.) 115 & #8201;pm

Covalent radius 133 & #8201;pm

Van der Waals radius 198 pm

Miscellaneous

Magnetic ordering nonmagnetic

Electrical resistivity (0 °C) 1.3×107 & #937;·m

Thermal conductivity (300 & #8201;K) 0.449

& #8201;W·m & #8722;1·K & #8722;1

Bulk modulus 7.7 & #8201;GPa

CAS registry number 7553-56-2

Selected isotopes

Main article: Isotopes of iodine iso NA half-life DM

DE (MeV) DP

127I 100% I is stable with 74 neutrons

129I syn 15.7×106y & #946;- 0.194 129Xe

131I syn 8.02070 d & #946;- 0.971 131Xe

References

For the record label, see Iodine Recordings.

Iodine (IPA: [ & #712;a & #618; & #601; & #716;da & #618;n],

/ & #712;a & #618; & #601; & #716;d & #618;n/, or

/ & #712;a & #618; & #601; & #716;di & #720;n/; from Greek:

iodes " violet " ), is a chemical element that has the

symbol I and atomic number 53. Chemically, iodine is

the least reactive of the halogens, and the most

electropositive halogen after astatine. Iodine is

primarily used in medicine, photography and dyes. It

is required in trace amounts by most living organisms.

As with all other halogens (members of Group VII in

the Periodic Table), iodine forms diatomic molecules,

and hence, has the molecular formula of I2.

Contents [hide]

1 Occurrence on earth

2 Uses

3 Isotopes

4 Notable characteristics

5 Sources

6 Descriptive chemistry

7 History

8 Notable inorganic iodine compounds

9 Stable iodine in biology

9.1 Human dietary intake

9.2 Iodine deficiency

9.3 Toxicity of iodine

10 Radioiodine and biology

10.1 Radioiodine and the thyroid

10.2 Radioiodine and the kidney

11 Non-hormone-related applications of iodine

12 Precautions for stable iodine

13 Clandestine use

14 See also

15 References

16 External links

[edit] Occurrence on earth

Iodine naturally occurs in the environment chiefly as

dissolved iodide in seawater, although it is also

found in some minerals and soils. The element may be

prepared in an ultrapure form through the reaction of

potassium iodide with copper(II) sulfate. There are

also a few other methods of isolating this element.

Although the element is actually quite rare, kelp and

certain other plants have the ability to concentrate

iodine, which helps introduce the element into the

food chain as well as keeping its cost down.

[edit] Uses

Iodine is used in pharmaceuticals, antiseptics,

medicine, food supplements, dyes, catalysts, halogen

lights, photography and water purifying.

[edit] Isotopes

There are 37 isotopes of iodine and only one, 127I, is

stable.

In many ways, 129I is similar to 36Cl. It is a soluble

halogen, fairly non-reactive, exists mainly as a

non-sorbing anion, and is produced by cosmogenic,

thermonuclear, and in-situ reactions. In hydrologic

studies, 129I concentrations are usually reported as

the ratio of 129I to total I (which is virtually all

127I). As is the case with 36Cl/Cl, 129I/I ratios in

nature are quite small, 10 & #8722;14 to 10 & #8722;10

(peak thermonuclear 129I/I during the 1960s and 1970s

reached about 10 & #8722;7). 129I differs from 36Cl in

that its half-life is longer (15.7 vs. 0.301 million

years), it is highly biophilic, and occurs in multiple

ionic forms (commonly, I & #8722; and IO3 & #8722;) which

have different chemical behaviors. This makes it

fairly easy for 129I to enter the biosphere as it

becomes incorporated into vegetation, soil, milk,

animal tissue, etc.

Excesses of stable 129Xe in meteorites have been shown

to result from decay of " primordial " 129I produced

newly by the supernovas which created the dust and gas

from which the solar system formed. 129I was the first

extinct radionuclide to be identified as present in

the early solar system. Its decay is the basis of the

I-Xe radiometric dating scheme, which covers the first

83 million years of solar system evolution.

Effects of various radioiodine isotopes in biology are

discussed below.

[edit] Notable characteristics

Iodine is a dark-gray/purple-black solid that sublimes

at standard temperatures into a purple-pink gas that

has an irritating odor. This halogen forms compounds

with many elements, but is less active than the other

members of its Group VII (halogens) and has some

metallic-like properties. Iodine dissolves easily in

chloroform, carbon tetrachloride, or carbon disulphide

to form purple solutions (It is only slightly soluble

in water, giving a yellow solution). The deep blue

color of starch-iodine complexes is produced only by

the free element.

Many students who have seen the classroom

demonstration where iodine crystals are gently heated

in a test tube come away with the impression that

liquid iodine cannot exist at atmospheric pressure.

This misconception arises because sublimation occurs

without the intermediacy of liquid. The truth is that

if iodine crystals are heated carefully to their

melting point of 113.7 °C, the crystals will fuse into

a liquid, which will be present under a dense blanket

of the vapour.

[edit] Sources

Iodine output in 2005Iodine is found in the mineral

Caliche, found in Chile, between the Andes and the

sea. It can also be found in some seaweeds as well as

extracted from seawater, however extracting Iodine

from the mineral is the only economical way to extract

the substance.[citation needed]

Extraction from seawater involves electrolysis, the

brine is first purified and acidified using sulphuric

acid and is then reacted with chlorine. An iodine

solution is produced but it is yet too dilute and has

to be concentrated. To do this air is blown into the

solution which causes the iodine to evaporate, then it

is passed into an absorbing tower containing acid

where sulfur dioxide is added to reduce the iodine,

the solution is then added to chlorine again to

concentrate the solution more, the final solution is

the iodine at a level of about 99%.[citation needed]

Another source is from kelp. This source was used in

the 18th and 19th centuries but is no longer

economically viable.

In 2005, Chile was the top producer of iodine with

almost two-thirds world share followed by Japan and

the USA reports the British Geological Survey.

[edit] Descriptive chemistry

Elemental iodine is poorly soluble in water, with one

gram dissolving in 3450 ml at 20 °C and 1280 ml at 50

°C. By contrast with chlorine, the formation of the

hypohalite ion (IO–) in neutral aqueous solutions of

iodine is negligible.

I2+ H2O & #8596; H+ + I– + HIO (K = 2.0×10-13) [1]

Solubility in water is greatly improved if the

solution contains dissolved iodides such as hydroiodic

acid, potassium iodide, or sodium iodide. Dissolved

bromides also improve water solubility of iodine.

Iodine is soluble in a number of organic solvents,

including ethanol (20.5 g/100 ml at 15 °C, 21.43 g/100

ml at 25 °C), diethyl ether (20.6 g/100 ml at 17 °C,

25.20 g/100 ml at 25 °C), chloroform, acetic acid,

glycerol, benzene (14.09 g/100 ml at 25 °C), carbon

tetrachloride (2.603 g/100 ml at 35 °C), and carbon

disulfide (16.47 g/100 ml at 25 °C)[2]. Aqueous and

ethanol solutions are brown. Solutions in chloroform,

carbon tetrachloride, and carbon disulfide are violet.

Elemental iodine can be prepared by oxidizing iodides

with chlorine:

2I– + Cl2 & #8594; I2 + 2Cl–

or with manganese dioxide in acid solution:[1]

2I– + 4H+ + MnO2 & #8594; I2 + 2H2O + Mn2+

Iodine is reduced to hydroiodic acid by hydrogen

sulfide:[3]

I2 + H2S & #8594; 2HI + S & #8595;

or by hydrazine:

2I2 + N2H4 & #8594; 4HI + N2

Iodine is oxidized to iodate by nitric acid:[4]

I2 + 10HNO3 & #8594; 2HIO3 + 10NO2 + 4H2O

or by chlorates:[4]

I2 + 2ClO3– & #8594; 2IO3– + Cl2

Iodine is converted in a two stage reaction to iodide

and iodate in solutions of alkali hydroxides (such as

sodium hydroxide):[1]

I2 + 2OH– & #8594; I– + IO– + H2O (K = 30)

3IO– & #8594; 2I– + IO3– (K = 1020)

[edit] History

Iodine was discovered by Bernard Courtois in 1811. He

was born to a manufacturer of saltpeter (a vital part

of gunpowder). At the time France was at war,

saltpeter, a component of gunpowder, was in great

demand. Saltpeter produced from French niter beds

required sodium carbonate, which could be isolated

from seaweed washed up on the coasts of Normandy and

Brittany. To isolate the sodium carbonate, seaweed was

burned and the ash then washed with water. The

remaining waste was destroyed by adding sulfuric acid.

One day Courtois added too much sulfuric acid and a

cloud of purple vapor rose. Courtois noted that the

vapor crystallized on cold surfaces making dark

crystals. Courtois suspected that this was a new

element but lacked the money to pursue his

observations.

However he gave samples to his friends,

Bernard Desormes (1777 - 1862) and Nicolas Clément

(1779 - 1841), to continue research. He also gave some

of the substance to ph Louis Gay-Lussac (1778 -

1850), a well-known chemist at that time, and to

André-Marie Ampère (1775 - 1836). On 29 November 1813,

Dersormes and Clément made public Courtois’ discovery.

They described the substance to a meeting of the

Imperial Institute of France. On December 6,

Gay-Lussac announced that the new substance was either

an element or a compound of oxygen. Ampère had given

some of his sample to Humphry Davy (1778 - 1829). Davy

did some experiments on the substance and noted its

similarity to chlorine. Davy sent a letter dated

December 10 to the Royal Society of London stating

that he had identified a new element. A large argument

erupted between Davy and Gay-Lussac over who

identified iodine first but both scientists

acknowledged Barnard Courtois as the first to isolate

the chemical element.

[edit] Notable inorganic iodine compounds

Ammonium iodide (NH4I)

Caesium iodide (CsI)

Copper(I) iodide (CuI)

Hydroiodic acid (HI)

Iodic acid (HIO3)

Iodine cyanide (ICN)

Iodine heptafluoride (IF7)

Iodine pentafluoride (IF5)

Lead(II) iodide (PbI2)

Lithium iodide (LiI)

Nitrogen triiodide (NI3)

Potassium iodide (KI)

Silver iodide (AgI)

Sodium iodide (NaI)

See also iodine compounds

[edit] Stable iodine in biology

Iodine is an essential trace element; its only known

roles in biology are as constituents of the thyroid

hormones, thyroxine (T4) and triiodothyronine (T3).

These are made from addition condensation products of

the amino acid tyrosine, and are stored prior to

release in a protein-like molecule called

thryroglobulin. T4 and T3 contain four and three atoms

of iodine per molecule, respectively. The thyroid

gland actively absorbs iodide ion from the blood to

make and release these hormones into the blood,

actions which are regulated by a second hormone TSH

from the pituitary. Thyroid hormones are

phylogenetically very old molecules which are

synthesized by most multicellular organisms, and which

even have some effect on unicellular organisms.

Thyroid hormones play a very basic role in biology,

acting on gene transcription to regulate the basal

metabolic rate. The total deficiency of thyroid

hormones can reduce basal metabolic rate up to 50%,

while in excessive production of thyroid hormones the

basal metabolic rate can be increased by 100%. T4 acts

largely as a precursor to T3, which is (with some

minor exceptions) the biologically active hormone.

[edit] Human dietary intake

The United States Food and Drug Administration

recommends 150 micrograms of iodine per day for both

men and women.[5] This is necessary for proper

production of thyroid hormone.[citation needed]

Natural sources of iodine include sea life, such as

kelp and certain seafood, as well as plants grown on

iodine-rich soil.[citation needed] Salt for human

consumption is often enriched with iodine and is

referred to as iodized salt.

[edit] Iodine deficiency

Main article: Iodine deficiency

In areas where there is little iodine in the

diet—typically remote inland areas and semi-arid

equatorial climates where no marine foods are

eaten—iodine deficiency gives rise to hypothyroidism,

symptoms of which are extreme fatigue, goitre, mental

slowing, depression, weight gain, and low basal body

temperatures.[citation needed]

Iodine deficiency is also the leading cause of

preventable mental retardation, an effect which

happens primarily when babies and small children are

made hypothyroid by lack of the element. The addition

of iodine to table salt has largely eliminated this

problem in the wealthier nations, but iodine

deficiency remains a serious public health problem in

the developing world.[citation needed]

[edit] Toxicity of iodine

Excess iodine has symptoms similar to those of iodine

deficiency. Commonly encountered symptoms are abnormal

growth of the thyroid gland and disorders in

functioning and growth of the organism as a whole.

Elemental iodine, I2, is a deadly poison if taken in

larger amounts; if 2-3 grams of it is consumed, it is

fatal to humans. Iodides are similar in toxicity to

bromides.

[edit] Radioiodine and biology

[edit] Radioiodine and the thyroid

The artificial radioisotope 131I (a beta emitter),

also known as radioiodine which has a half-life of

8.0207 days, has been used in treating cancer and

other pathologies of the thyroid glands. 123I is the

radioisotope most often used in nuclear imaging of the

kidney and thyroid as well as thyroid uptake scans

(used for the evaluation of Grave's disease). The most

common compounds of iodine are the iodides of sodium

and potassium (KI) and the iodates (KIO3).

129I (half-life 15.7 million years) is a product of

130Xe spallation in the atmosphere and uranium and

plutonium fission, both in subsurface rocks and

nuclear reactors. Nuclear processes, in particular

nuclear fuel reprocessing and atmospheric nuclear

weapons tests have now swamped the natural signal for

this isotope. 129I was used in rainwater studies

following the Chernobyl accident. It also has been

used as a ground-water tracer and as an indicator of

nuclear waste dispersion into the natural environment.

If humans are exposed to radioactive iodine, the

thyroid gland will absorb it as if it were

non-radioactive iodine, leading to elevated chances of

thyroid cancer. Isotopes with shorter half-lives such

as 131I present a greater risk than those with longer

half-lives since they generate more radiation per unit

of time. Taking large amounts of regular iodine will

saturate the thyroid and prevent uptake. Iodine pills

are sometimes distributed to persons living close to

nuclear establishments, for use in case of accidents

that could lead to releases of radioactive iodine.

Iodine-123 and iodine-125 are used in medicine as

tracers for imaging and evaluating the function of the

thyroid.

Iodine-131 is used in medicine for treatment of

thyroid cancer and Grave's disease.

Uncombined (elemental) iodine is mildly toxic to all

living things.

Potassium iodide (KI tablets, or " SSKI " =

" Super-Saturated KI " liquid drops) can be given to

people in a nuclear disaster area when fission has

taken place, to flush out the radioactive iodine-131

fission product. The half-life of iodine-131 is only

eight days, so the treatment would need to continue

only a couple of weeks. In cases of leakage of certain

nuclear materials without fission, or certain types of

dirty bomb made with other than radioiodine, this

precaution would be of no avail.

[edit] Radioiodine and the kidney

In the 1970s imaging techniques were developed in

California to utilize radioiodine in diagnostics for

renal hypertension.

[edit] Non-hormone-related applications of iodine

Tincture of iodine (5% elemental iodine in

water/ethanol base) is an essential component of any

emergency survival kit, used both to disinfect wounds

and to sanitize surface water for drinking (3 drops

per litre, let stand for 30 minutes). Alcohol-free

iodine solutions such as Lugol's iodine, as well as

other iodophor type antiseptics, are also available as

effective elemental iodine sources for this purpose.

Iodine compounds are important in the field of organic

chemistry

Iodine, as a heavy element, is quite radio-opaque.

Organic compounds of a certain type (typically

iodine-substituted benzene derivatives) are thus used

in medicine as X-ray radiocontrast agents for

intravenous injection. This is often in conjunction

with advanced X-ray techniques such as angiography and

CT scanning

Silver iodide is used in photography.

Tungsten iodide is used to stabilize the filaments in

light bulbs.

[edit] Precautions for stable iodine

Direct contact with skin can cause lesions, so it

should be handled with care. Iodine vapor is very

irritating to the eye and to mucous membranes.

Concentration of iodine in the air should not exceed 1

mg/m³ (eight-hour time-weighted average). When mixed

with ammonia, it can form nitrogen triiodide which is

extremely sensitive and can explode unexpectedly.

[edit] Clandestine use

In the United States, the Drug Enforcement Agency

(DEA) regards iodine and compounds containing iodine

(ionic iodides, iodoform, ethyl iodide, and so on) as

reagents useful for the clandestine manufacture of

methamphetamine. Persons who attempt to purchase

significant quantities of such chemicals without

establishing a legitimate use are likely to find

themselves the target of a DEA investigation. Persons

selling such compounds without doing due diligence to

establish that the materials are not being diverted to

clandestine use may be subject to stiff penalties,

such as expensive fines or even imprisonment.[6][7]

________________________________________________________________________________\

____

Choose the right car based on your needs. Check out Autos new Car Finder

tool.

http://autos./carfinder/