Iodine

What is Iodine?

Iodine is a chemical element with the atomic symbol I and atomic number 53. One of the most massive stable halogens, it is a semi-lustrous, non-metallic solid at room temperature that melts to a deep violet liquid at 114 °C (237 °F) and boils to a violet gas at 184 °C (363 °F) under standard conditions.

The element was named after the Ancient Greek term Iώδης, which means “violet-colored,” by Joseph Louis Gay-Lussac two years after it was discovered by the French chemist Bernard Courtois in 1811.

Iodine can be found in numerous oxidation states, such as iodate (IO⁻₃), iodide (I ⁻), and the different periodate anions.

As the sixty-first most abundant element, it is the least abundant of the stable halogens. Iodine is the most densely packed essential mineral ingredient and is necessary for the production of thyroid hormones.

Approximately two billion individuals worldwide suffer from iodine deficiency, which is the main avoidable cause of intellectual disability.

These days, Japan and Chile are the world’s leading iodine producers. As a non-toxic radiocontrast substance, it has also gained popularity because of its high atomic number and ease of attachment to organic molecules.

Thyroid cancer can also be treated using radioactive isotopes of iodine due to the specificity of their absorption by the human body.

Additionally, acetic acid and other polymers are produced industrially using iodine as a catalyst. It is included in the List of Essential Medicines by the World Health Organization.

History

French chemist Bernard Courtois, who was born into a family that produced saltpeter, an ingredient needed to make gunpowder, made the discovery of iodine in 1811.

France had a high demand for saltpeter during the Napoleonic Wars. Sodium carbonate, which could be separated from seaweed gathered along the coastlines of Normandy and Brittany, was needed to manufacture saltpeter from French niter beds.

Seaweed was burned, and the ash was then cleaned with water to separate the sodium carbonate. Sulfuric acid was added to eliminate the leftover garbage.

One time, Courtois put too much sulfuric acid, and a plume of purple vapor erupted. He observed that the vapor formed dark crystals when it crystallized on cold surfaces. Courtois lacked the resources to investigate his suspicions that this material was novel.

To facilitate further investigation, Courtois provided samples to his pals Nicolas Clément (1779–1841) and Charles Bernard Desormes (1777–1838).

Additionally, he gave some of the material to André-Marie Ampère (1775–1836), a physicist, and Joseph Louis Gay-Lussac (1778–1850), a chemist.

Desormes and Clément revealed Courtois’ discovery to the world on November 29, 1813. They presented the details to the Imperial Institute of France meeting.

Gay-Lussac said on December 6 that the new material was either an element or an oxygen compound.

Gay-Lussac proposed the term “iode” based on the color of iodine vapor, derived from the Ancient Greek word ἰoειδής (ioeidēs, “violet”).

Ampère had provided He gave a portion of his sample to the English chemist Humphry Davy (1778–1829), who investigated the material and discovered that it resembled chlorine.

Davy informed the Royal Society of London in a letter dated December 10 that he had discovered a new element.

Though there were disagreements between Davy and Gay-Lussac on who discovered iodine initially, both scientists agreed that Courtois was the first to isolate the element.

Casimir Joseph Davaine (1812–1882), a French medical researcher, discovered iodine’s antibacterial properties in 1873.

Antonio Grossich (1849–1926), a surgeon from Istriana, was one of the pioneers in the application of operating room sterilization. He invented the tincture of iodine in 1908 to quickly sterilize human skin during surgery. Iodine was frequently represented by the sign J in early periodic tables, after Jod, its German name.

Properties

As the heaviest stable member of its group, iodine is the fourth halogen and belongs to group 17 of the periodic table, which is positioned below fluorine, chlorine, and bromine.

(The radioactive astatine and tennessine, the fifth and sixth halogens, are not frequently investigated because they are expensive and difficult to obtain in significant quantities, but relativistic effects seem to give them certain odd features for the group.)

The electron configuration of iodine is  [Kr]4d¹⁰5s²5p⁵, with its valence consisting of seven electrons in the fifth and outermost shell. electrons.

Although it is an oxidizing agent like the other halogens because it is one electron away from completing an octet and reacts with a variety of elements to complete its outer shell, it is the least potent oxidizing agent among them.

The stable halogens due to their lowest electronegativity, 2.66 on the Pauling scale (compare fluorine, chlorine, and astatine, with an electronegativity of 2.2, continue the trend, with bromine at 3.98, 3.16, and 2.96, respectively.

Therefore, two iodine atoms share a pair of electrons to form a stable octet for themselves; at high temperatures, these diatomic molecules reversibly dissociate a pair of iodine atoms.

This is how elemental iodine becomes diatomic molecules with chemical formula I2. In a similar vein, of all the stable halogens, the iodide anion, I⁻, is the strongest reducing agent since it can be oxidized back to diatomic I₂ the easiest.

(Actine is even more unstable, starting as At⁻ and easily oxidizing to At0 or At⁺.)As the group descends, the halogens’ colors darken: iodine is violet, bromine is reddish-brown, fluorine is extremely pale yellow, and chlorine is greenish-yellow.

One gram of elemental iodine dissolves in 3450 milliliters at 20°C and 1280 milliliters at 50°C in water. Potassium iodide can be added to boost solubility by forming triiodide ions among other polyiodides.

A higher solubility is offered by nonpolar solvents like carbon tetrachloride and hexane. Nonpolar solutions are violet, the hue of iodine vapor; polar solutions, like aqueous solutions, are brown, reflecting these solvents’ function as Lewis bases.

Iodine changes color when dissolved in polar liquids because charge-transfer compounds are formed.

Iodine is deep brown in alcohols and amines, solvents that create charge-transfer adducts, while it appears violet when dissolved in carbon tetrachloride and saturated hydrocarbons.

Since iodine has the largest and most easily polarized electron cloud among the halogens, its molecules have the strongest van der Waals interactions.

As a result, iodine has the highest melting and boiling points among the halogens, following the increasing trend down the group. Comparably, iodine is the least volatile of the halogens, but it nevertheless exhibits purple light in solid form. vapor.

Owing to this characteristic, iodine is frequently utilized to illustrate the direct sublimation of solids into gases, leading to the misperception that it does not melt at room temperature.

Iodine has the lowest first ionization energy, lowest electron affinity, lowest electronegativity, and lowest reactivity of all the halogens due to its greatest atomic radius. Of all the halogens, the interhalogen bond in diiodine is the weakest.

Therefore, at 575 °C, 1% of a gaseous iodine sample at atmospheric pressure dissociates into iodine atoms. In order for fluorine, chlorine, and bromine to dissociate to a comparable degree, temperatures higher than 750 °C are necessary.

Compared to the equivalent bonds to the lighter halogens, the majority of bonds to iodine are weaker. I₂ molecules with an I–I bond length of 266.6 pm make up gaseous iodine.

One of the longest-known single bonds is the I–I bond. Well, it is in solid orthorhombic crystalline iodine, which shares a crystal structure with bromine and chlorine, even longer (271.5 pm).

(Xenon, iodine’s neighbor, holds the record; the Xe–Xe bond length is 308.71 pm.) Because of this, there are substantial electronic interactions between each atom and its two next-nearest neighbors within the iodine molecule.

These interactions give bulk iodine its shiny look and semiconducting qualities. With a band gap of 1.3 eV (125 kJ/mol), iodine is a two-dimensional semiconductor that is an insulator perpendicular to its crystalline layers and a semiconductor in the plane of its layers.

Isotopes

Iodine-127 is the only one of the thirty-seven known isotopes of iodine that can be found in nature.

The others are too short in half-lives to be primordial, and they are radioactive. Iodine is therefore monoisotopic and mononuclidic, and since it is a constant of nature, its atomic weight is known with extreme precision.

Iodine-129, which decays by beta decay to stable xenon-129, has the longest half-life of any radioactive isotope of iodine, at 15.7 million years.

Iodine-129 and iodine-127 were partially created before the Solar System, but they have since entirely decayed, making it an extinct radionuclide that is nonetheless helpful for dating the early Solar System or extremely old groundwaters. its ability to move around in the surroundings.

Its offspring xenon-129 excess can be used to determine its previous presence. Since iodine-129 is also a cosmogenic nuclide that was created by cosmic ray spallation of atmospheric xenon, traces of it are still present today and account for 10⁻¹⁴ to 10⁻¹⁰ of all iodine on Earth.

The longest half-life of any fission product also results from open-air nuclear testing and is not dangerous.

Even during the height of thermonuclear testing in the 1960s and 1970s, iodine-129 accounted for only approximately 10^^-7 of the total iodine found on Earth.

In Mössbauer spectroscopy, excited states of iodine-127 and iodine-129 are frequently employed. The half-lives of the other iodine radioisotopes are only a few days, significantly shorter.

A few of them have therapeutic uses related to the thyroid gland, which is where the body stores and concentrates the iodine that enters.

With a half-life of thirteen hours, iodine-123 decays by electron capture to tellurium-123, which releases gamma radiation.

It is utilized in nuclear medicine imaging procedures such as X-ray computed tomography (X-ray CT) scans and single photon emission computed tomography (SPECT).

Having a half-life of fifty-nine days, iodine-125 decays to tellurium-125 through electron capture, releasing low-energy gamma radiation.

Being the second-longest-lived iodine radioisotope, it is used in nuclear medicine imaging, biological assays, and radiation therapy (brachytherapy) to treat various conditions like prostate cancer uveal brain tumors, and melanomas.

Ultimately, iodine-131, which has an eight-day half-life, beta decays to an excited state of stable xenon-131, which then releases gamma radiation to return to the ground state.

Since it is a frequent fission product, radioactive fallout contains large amounts of it. It will then build up in the thyroid and might be absorbed through tainted food.

It may harm the thyroid as it decomposes. The main danger of prolonged exposure to high iodine-131 concentrations is the possibility of developing radiogenic thyroid cancer in later life. Thyroiditis and non-cancerous growths are potential additional concerns.

Typically, the thyroid gland is saturated with stable iodine-127 in the form of potassium iodide tablets, which should be taken daily for optimal prophylaxis, to guard against the harmful effects of iodine-131.

However, for just this reason—when tissue damage is sought following iodine uptake by the tissue— iodine-131 may also be used medicinally in radiation therapy. Another radioactive tracer that is utilized is iodine-131.

Chemistry and Compounds

Although iodine is very reactive, it is not as reactive as the other halogens. Chlorine gas, on the other hand, will halogenate CO, NO, and SO2 (becoming phosgene, nitrosyl chloride, and sulfuryl chloride, respectively).

Iodine, on the other hand, will not do this. Moreover, iodination of metals, as opposed to chlorination or bromination, usually results in lower oxidation states; for instance, rhenium metal interacts with chlorine to generate rhenium hexachloride.

However, only rhenium pentabromide can be produced with bromine, and only rhenium tetraiodide can be produced with iodine.

Conversely, iodine has more significant cationic chemistry and its higher oxidation states are much more stable than those of bromine and chlorine, for example in iodine heptafluoride, because it has the lowest ionization energy of all the halogens and is the most easily oxidized of them.

Charge-Transfer Complexes

Bright violet solutions are produced when the iodine molecule, I₂, dissolves in CCl₄ and aliphatic hydrocarbons.

The absorption band maximum in these solvents is attributed to a π* to σ* transition and lies in the 520–540 nm area.

In these solvents, I₂ interacts with Lewis bases to generate charge-transfer complexes (adducts), which cause a blue shift in the I₂ peak and the emergence of a new peak (230–330 nm).

Hydrogen Iodide

Hydrogen iodide (HI), the simplest form of iodine, is a chemical. It is a colorless gas that produces water and iodine when it combines with oxygen.

Like the other hydrogen halides, it is helpful in laboratory iodination reactions but has no industrial applications on a wide scale.

It is often produced commercially by reacting iodine with hydrazine or hydrogen sulfide:

2 I₂ + N₂H₄ H₂O⟶ 4 HI + N₂

Like all hydrogen halides, with the exception of hydrogen fluoride, it is a colorless gas at ambient temperature because hydrogen is unable to establish strong hydrogen bonds with the massive and only slightly electronegative iodine atom. It boils at -35.1 °C and melts at -51.0 °C.

It is an endothermic molecule that, in the absence of a catalyst, can exothermically dissociate at room temperature.

Specifically, the reaction between hydrogen and iodine at ambient temperature that yields hydrogen iodide is not completed. At 295 kJ/mol, the H–I bond dissociation energy is also the lowest of the hydrogen halides.

Since hydrogen is unable to make strong hydrogen bonds with the big and only slightly electronegative iodine atom, it is a colorless gas at room temperature, just like all other hydrogen halides except hydrogen fluoride.

At -51.0 °C, it melts, and at -35.1 °C, it boils. This chemical is endothermic and can dissociate at room temperature exothermically, but the process requires a catalyst and proceeds extremely slowly.

At room temperature, the reaction between hydrogen and iodine that yields hydrogen iodide does not fully proceed. Similarly, the H–I bond dissociation energy, at 295 kJ/mol, is the lowest among the hydrogen halides.

Strong acid hydroiodic acid is the term used to describe aqueous hydrogen iodide. One liter of water will dissolve 425 liters of hydrogen iodide, and in a saturated solution, there are just four water molecules for every hydrogen iodide molecule.

This indicates that hydrogen iodide is incredibly soluble in water. Commercial hydroiodic acid that is referred to as “concentrated” typically has a mass percentage of 48–57% HI; at 56.7 g HI per 100g solution, an azeotrope forms with a boiling temperature of 126.7 °C.

Therefore, the evaporation of water cannot concentrate hydroiodic acid beyond this limit. Hydroiodic acid, as opposed to gaseous hydrogen iodide, is primarily used in industry to produce acetic acid by the Cativa process.

In contrast to hydrogen fluoride, anhydrous liquid hydrogen iodide presents challenges when used as a solvent due to its low boiling point, narrow liquid range, low permittivity, and inability to dissociate significantly into H₂I⁺ and HI⁻₂ ions.

The latter are less stable than fluoride ions (HF⁻₂) because of the weak hydrogen bonding between iodine and hydrogen, although they can still be isolated in their salt forms with large and weakly polarizing cations like Cs⁺ and NR⁺₄ (R = Me, Et, Buⁿ) in them.

Anhydrous hydrogen iodide is a weak solvent; it can only dissolve very low lattice energy salts, like tetraalkylammonium halides, or tiny molecules, like phenol and nitrosyl chloride.

Other Binary Iodine Compounds

Nearly every element on the periodic table up to einsteinium (EsI₃ is known), except the noble gases, is known to form binary compounds with iodine. Nitrogen triiodide was solely recognized as an ammonia adduct until 1990.

It was discovered that ammonia-free NI₃ is isolable at -196 °C but spontaneously decomposes at 0 °C.

Neutral sulfur and selenium iodides that are stable at room temperature also do not exist for thermodynamic reasons relating to the electronegativity of the elements; nonetheless, S₂I₂ and SI₂ are stable up to 183 and 9K, respectively.

No unequivocally identifiable neutral binary selenium iodide has been found as of 2022 (at any temperature).

Crystallographic characterization and preparation of sulfur- and selenium-iodine polyatomic cations (e.g., [S₂I₄²⁺][AsF₆^–]₂ and [Se₂I₄²⁺][Sb₂F₁₁^–]₂) have been carried out.

High oxidation states are hard to attain in binary iodides due to the huge size of the iodide anion and the weak oxidizing power of iodine; the maximum known is found in the pentaiodides of niobium, tantalum, and protactinium.

Hydroiodic acid with an element or its oxide, hydroxide, or carbonate can react to produce iodides.

The resulting gas is subsequently dehydrated using a combination of low pressure and gently elevated temperatures, or by using anhydrous hydrogen iodide gas.

When the iodide product is stable for hydrolysis, these techniques function best. The element can also be synthesized via high-temperature oxidative iodination.

High-temperature iodination of a metal oxide or other halide by iodine, a volatile metal halide, carbon tetraiodide, or an organic iodide, with iodine or hydrogen iodide.

For instance, at 230 °C, molybdenum(IV) oxide and aluminum(III) iodide react to produce molybdenum(II) iodide. The following is an example of halogen exchange: at 400 °C, tantalum(V) chloride reacts with excess aluminum(III) iodide to produce tantalum(V) iodide.

3TaI5 + 5AlCl3}}}”>
Lower iodides can be created by reducing the higher iodide with hydrogen or a metal, for instance, or by thermal breakdown or disproportionation.

The majority of metal iodides are ionic when the metal is in low oxidation states (+1 to +3). Covalent molecular iodides are typically formed by nonmetals and metals in high oxidation states (above +3).

For metals in oxidation state +3, ionic and covalent iodides are known (for example, aluminum iodide is not primarily ionic, but scandium iodide is).

Among the halides MX_n of the same element, ionic iodides MI_n typically have the lowest melting and boiling temperatures due to the big iodide anion’s weaker electrostatic forces of attraction between the cations and anions.

Conversely, covalent iodides typically contain the greatest melting and boiling temperatures among the halides of the same element because iodine may contribute the most to van der Waals forces due to its high electron count and ability to polarize the most among halogens.

In intermediate iodides, of course, there are plenty of instances where one trend gives way to the other.

Comparably, among the ionic halides of that element, the solubilities of primarily ionic iodides (like potassium and calcium) in water are the highest, whilst those of covalent iodides (like silver) are the lowest.

Specifically, silver iodide is highly soluble in water, and its precipitation is frequently employed as a qualitative indicator of iodine.

Iodine Halides

Iodine is not an exception to the rule that the halogens create several binary, diamagnetic interhalogen compounds with stoichiometries XY, XY₃, XY₅, and XY₇ (where X is heavier than Y).

Iodine creates pentafluoride, trichloride, trifluoride, and, incredibly among the halogens, heptafluoride.

It also makes all three conceivable diatomic interhalogens. A wide range of cationic and anionic derivatives are also characterized; examples include the dark brown or purplish black compounds of I₂Cl⁺ and the wine-red or bright orange compounds of ICl⁺₂.

In addition to these, other pseudohalides are also recognized, including iodine azide (IN3), cyanogen iodide (ICN), and iodine thiocyanate (ISCN).

It is impossible to obtain pure iodine monofluoride (IF) because it is unstable at ambient temperature and easily and irreversibly disproportionates to iodine and iodine pentafluoride.

The process of synthesizing it involves reacting iodine with fluorine gas in trichlorofluoromethane at -45°C, iodine trifluoride in trichlorofluoromethane at -78°C, or silver(I) fluoride at zer0°C.

Conversely, iodine monochloride (ICl) and iodine monobromide (IBr) have a modest level of stability.

Shortly after the discoveries of iodine and chlorine, Joseph Louis Gay-Lussac and Humphry Davy separately discovered the former, a volatile reddish-brown molecule that mimics the intermediate halogen bromine so well that Justus von Liebig was duped into thinking it was bromine.

(that he had discovered) for monochloroiodine. Iodine can be made into iodine monochloride and iodine monobromide by simply reacting it with either bromine or chlorine at room temperature, and then purifying it via fractional crystallization.

They are both highly reactive and can even attack gold and platinum, but not tungsten, boron, carbon, cadmium, lead, zirconium, niobium, molybdenum, or zinc.

The circumstances determine how they react with organic molecules. Since iodine and chlorine are formed during the homolytic fission of iodine chloride, and iodine is more reactive than chlorine, iodine chloride vapor tends to chlorinate salicylic acid and phenol.

But when iodine chloride is added to a carbon tetrachloride solution, the primary reaction is iodination because the I-Cl bond is now heterolytically fissioned.

Transpires, and as an electrophile, I⁺ targets phenol. Iodine monobromide, on the other hand, tends to brominate phenol even in carbon tetrachloride solution because bromine is more reactive than iodine and it tends to dissociate into its constituents in solution.

Iodine monochloride and iodine monobromide, when liquid, separate into I₂X⁺ and IX⁻₂ ions (X = Cl, Br); as a result, they are useful electrical conductors and can be employed as ionizing solvents.

Iodine trifluoride (IF₃) is a yellow solid that is unstable and breaks down above -28 °C. Thus, little is known about it.

It is challenging to make because xenon difluoride must react at a low temperature with fluorine gas, which has the tendency to oxidize iodine all the way to pentafluoride.

Bright yellow in solid form, iodine trichloride is created by reacting iodine with liquid chlorine at a temperature of -80°C.

It should be handled carefully during purification as it can act as a strong chlorinating agent and readily dissociate into iodine monochloride and chlorine.

The electrical conductivity of liquid iodine trichloride may suggest dissociation to ICl⁺₂ and ICl⁻ ₄ ions.

The most thermodynamically stable form of iodine fluoride is iodine pentafluoride (IF₅), a colorless, volatile liquid that is created by combining iodine with fluorine gas at ambient temperature.

Although it is a fluorinating agent, it is not strong enough to be kept in glass containers. Once more, the liquid state exhibits a small amount of electrical conductivity due to dissociation to IF⁺ ₄ and IF⁻₆.

Next only to chlorine trifluoride, chlorine pentafluoride, and bromine pentafluoride among the interhalogens in terms of potency, is the pentagonal bipyramidal iodine heptafluoride (IF₇).

Which reacts with nearly all elements even at low temperatures, fluorinates Pyrex glass to form iodine(VII) oxyfluoride (IOF₅), and ignites carbon monoxide.

Iodine Oxides and Oxoacids

Because of the strong I–O bonds that arise from the significant electronegativity difference between iodine and oxygen, iodine oxides are the most stable of all the halogen oxides and have been for the longest.

Since its discovery in 1813 by Gay-Lussac and Davy, the stable, white, hygroscopic iodine pentoxide (I₂O₅) has been understood to exist.

The most straightforward way to make it is to dehydrate iodic acid (HIO₃), of which it is the anhydride.

It is a helpful reagent for figuring out carbon monoxide concentration since it will rapidly oxidize carbon monoxide fully to carbon dioxide at ambient temperature.

Moreover, it oxidizes hydrogen sulfide, ethylene, and nitrogen oxide. Sulfur trioxide and its reaction. Concentrated sulfuric acid reduces peroxydisulfuryl difluoride (S₂O₆F₂) to iodosyl salts involving [IO]⁺, which are salts of the iodyl cation, [IO₂]⁺.

Iodine pentafluoride, which also reacts with iodine pentoxide to produce iodine(V) oxyfluoride, or IOF₃, is the result of fluorinating it with fluorine, bromine trifluoride, sulfur tetrafluoride, or chloryl fluoride.

Some less stable oxides are known, such as I₄O₉ and I₂O₄. Although their structures are unknown, it is feasible to assume that they are  I^III(I^VO₃)₃  and [IO]⁺[IO₃]⁻, respectively.

The four oxoacids—periodic acid (HIO₄ or H5IO₆), iodic acid (HIO₃), hypoiodous acid (HIO), and iodous acid (HIO₂)—are more significant. The following processes take place when iodine dissolves in an aqueous solution:

It is unstable to disproportionation, hypoiodous acid. Thus, disproportionately quickly, the hypoiodite ions formed to produce iodide and iodate:
3 IO⁻ ⇌ 2 I⁻ + IO⁻₃K = 10²⁰

Even less stable are iodous acid and iodite, which are only ever present as a transient step in the oxidation of iodide to iodate.

The most significant of these compounds are iodates, which are produced by oxidizing iodine with chlorates or alkali metal iodides with oxygen at 600 °C and high pressure.

Iodates are stable against disproportionation in both acidic and alkaline solutions, in contrast to chlorates, which are disproportionate very slowly to form chloride and perchlorate. Most metal salts can be made from these.

The simplest methods for producing iodic acid include fuming nitric acid or electrolyzing an aqueous iodine solution. Among the halates, iodate responds the fastest but has the least oxidizing power.

There are numerous known periodates, including the anticipated tetrahedral IO−4, as well as square-pyramidal IO₃⁻₅, octahedral orthoperiodate IO⁵⁻₆, [IO₃(OH)₃]²⁻, [I₂O₈(OH₂)]⁴⁻.

Typically, they are produced by either using chlorine gas or electrochemically oxidizing alkaline sodium iodate (with lead(IV) oxide as the anode):

IO⁻₃ + 6 OH⁻ → IO⁵⁻₆ + 3 H₂O + 2 e⁻

IO⁻₃ + 6 OH⁻ + Cl₂ → IO⁵⁻₆ + 2 Cl⁻ + 3 H₂O

They are strong oxidizers both thermodynamically and kinetically, converting Mn²⁺ to MnO⁻₄ fast and cleaving glycols, α-diketones, α-ketols, α- amino alcohols, and α-diamines.

Orthoperiodate’s extremely high negative charge of −5 specifically stabilizes high oxidation states in metals. Stable at 100 °C in a vacuum, orthoperiodic acid (H5IO₆) dehydrates to metaperiodic acid (HIO₄).

Trying harder yields iodine pentoxide and oxygen instead of the nonexistent iodine heptoxide (I₂O₇).

Sulfuric acid can protonate periodic acid, resulting in the I(OH)⁺₆ cation, which is isoelectronic to Te(OH)₆ and Sb(OH)⁻₆, and salts including bisulfate and sulfate.

Polyiodine Compounds

I⁺₂ cations and a vivid blue paramagnetic solution are produced when iodine dissolves in strong acids, such as fuming sulfuric acid. Antimony pentafluoride can be used to oxidize iodine to produce a solid salt of the diiodine
2 I₂ + 5 SbF₅ SO₂⟶20 °C 2 I₂Sb₂F₁₁ + SbF₃

 I₂Sb₂F₁₁, the blue tantalum counterpart, is also known, and the salt I₂Ta₂F₁₁ is dark blue. I–I bond length in I⁺₂ is only 256 pm, compared to 267 pm in I₂ because the latter’s missing electron was taken out of an antibonding orbital, which made the bond stronger and shorter.

Deep-blue I⁺₂ reversibly dimerizes in fluorosulfuric acid solution below −60 °C to generate red rectangular diamagnetic I2⁺₄.

Some polyiodine cations, such as bent dark-brown or black I⁺₃ and centrosymmetric C_2h green or black I⁺₅, which are present in the AsF− 6 and AlCl− 4 salts, are not as well characterized.

In an aqueous solution, linear triiodide, I⁻₃, is the only significant polyiodide anion. The process of its creation explains why adding potassium iodide solution may make iodine more soluble in water:

I₂ + I⁻ ⇌ I⁻₃ (K_eq = ~700 at 20 °C)
When solutions containing iodine and iodide crystallize, many additional polyiodides, such as  I⁻
₅, I⁻₉, I²⁻₄, and I²⁻₈, and I₂⁻₈, may be obtained. These solutions may also allow for the isolation of their salts, which have massive, weakly polarizing cations like Cs⁺.

Organoiodine Compounds

Organoiodine compounds have played a pivotal role in the advancement of organic synthesis, exemplified by their application in the Wurtz coupling reaction, Williamson ether synthesis, Hofmann elimination of amines, and Grignard reagents.

Formally, these compounds can be considered organic derivatives of the iodide anion. The carbon–iodine bond is a typical functional group that is a part of core organic chemistry.

Alkyl iodides, the most basic type of organoiodine compounds, can be created by reacting alcohols with phosphorus triiodide; these can subsequently be utilized in Grignard reagent preparation or nucleophilic substitution processes.

The tiny difference in electronegativity between carbon (2.55) and iodine (2.66) makes the C–I bond the weakest of all the carbon–halogen bonds.

Iodide is, therefore, the finest leaving group among the halogens; in fact, so many organoiodines Due to the facile production and breakage of the C–I bond, compounds are frequently utilized in organic synthesis, however, they eventually turn yellow when stored due to breakdown into elemental iodine.

The high atomic weight of iodine in them also makes them substantially denser than the other organohalogen compounds.

Certain oxidizing agents that are organic, such as iodanes, have iodine in an oxidation state higher than −1, like 2-iodobenzene dichloride (PhICl₂), which is used to selectively chlorinate alkenes and alkynes, and iodoxybenzoic acid, a typical reagent for the oxidation of alcohols to aldehydes.

The so-called “iodoform test,” which produces iodoform (CHI₃) by exhaustively iodinating a methyl ketone (or another chemical that can be oxidized to a methyl ketone), is one of the more well-known applications of organoiodine compounds.

It works as follows: Since iodine is expensive and organoiodine compounds are stronger alkylating agents, there are certain disadvantages to employing organoiodine compounds as opposed to organochlorine or organobromine compounds, including the higher cost and toxicity of the iodine derivatives.

Iodoacetamide and iodoacetic acid, for instance, denature proteins by inhibiting the formation of disulfide bonds and permanently alkylating cysteine residues.

The fact that iodide is a better-leaving group than chloride or bromide significantly complicates the halogen exchange process used in the Finkelstein reaction to make iodoalkanes.

Nevertheless, the difference is negligible enough to allow the reaction to be fully driven by employing a significant excess of the halide salt or by taking advantage of the halide salts’ varying solubility.

In the traditional Inkelstein process, an alkyl chloride or alkyl bromide is treated with a sodium iodide solution in acetone to produce an alkyl iodide.

While sodium chloride and sodium bromide are not soluble in acetone, sodium iodide is. The precipitation of the insoluble salt causes mass action to move the reaction toward products.

Occurrence and Production

Among the stable halogens, iodine is the least plentiful, making up only 0.46 parts per million of the crustal rocks of Earth (flux is 544 ppm, chlorine is 126 ppm, and bromine is 2.5 ppm).

It is the 61st most abundant element out of the 83 elements (elements 1-42, 44-60, 62-83, 90, and 92) that are found in substantial quantities.

The majority of deposits that are concentrated enough for profitable extraction are made up of iodate minerals rather than iodide minerals, which are rare. Lautarite, Ca(IO₃)₂, and dietzeite, 7Ca(IO₃)₂·8CaCrO₄, are two examples.

These are the minerals that can be found in trace amounts in Chilean caliche. whose sodium nitrate is its principal product.

They can have a maximum of 1% and a minimum of 0.02% iodine by mass overall. Sodium bisulfite is used to remove sodium iodate from the caliche and convert it to iodide.

After that, this solution reacts with recently extracted iodate to produce proportionate iodine, which can be removed by filtering.

Brine became a comparable source in the late 20th century, although caliche remained the primary source of iodine in the 19th century and is still significant today, displacing kelp, which is no longer a viable source.

The two largest of these sources are the American Anadarko Basin gas field in northwest Oklahoma and the Japanese Minami Kanto gas field east of Tokyo.

At the source’s depth, the brine is hotter than 60 °C. Sulfuric acid is used to purify and acidify the brine, and chlorine is then used to oxidize the iodide in the brine to iodine. It produces an iodine solution, but it needs to be concentrated because it is diluted.

The iodine is evaporated by blowing air into the solution, and then it is fed into an absorbent tower where sulfur dioxide breaks down the iodine.

Iodine is precipitated through a reaction between hydrogen iodide (HI) and chlorine. Iodine is packed following purification and filtering.

2 HI + Cl₂ → I₂↑ + 2 HCl
I₂ + 2 H₂O + SO₂ → 2 HI + H₂SO₄
2 HI + Cl₂ → I₂↓ + 2 HCl

These sources guarantee that the two biggest iodine producers in the world today are Chile and Japan.

An alternative would be to treat the brine with silver nitrate, which would cause the iodine to precipitate out as silver iodine.

This silver iodide would then react with the iron to generate metallic silver and an iron(II) iodide solution. Chlorine displacement can then be used to release the iodine.

Applications

Around half of the iodine that is produced is used to create different organoiodine compounds, 15% is left as the element in its pure form, 15% is used to create potassium iodide, and the remaining 15% is used to create other inorganic iodine compounds.

Catalysts, stabilizers, colorants, pigments, pharmaceuticals, photography, sanitation (using iodine tincture), and cloud seeding are a few of the main applications for iodine compounds; other uses include smog inhibition and different applications in analytical chemistry.

Chemical Analysis

Iodometry is one application of quantitative volumetric analysis that frequently uses the iodide and iodate anions.

Iodine and starch combine to generate a blue complex, which is frequently used as an indicator in iodometry and to test for iodine or starch.

To identify fake banknotes printed on starch-containing paper, the iodine test for starch is still utilized. The mass of iodine in grams that is consumed by 100 grams of a chemical substance, usually fats or oils, is known as the iodine value.

It is common practice to utilize iodine values to estimate the degree of unsaturation in fatty acids. Double bonds, which result from this unsaturation, react with iodine compounds.

Nessler’s reagent, or potassium tetraiodomercurate(II), K₂HgI₄, is another name for it. It is frequently employed as an ammonia-sensitive spot test.

Similar to this, alkaloids are detected using Mayer’s reagent, which is a potassium tetraiodomercurate(II) solution employed as a precipitating reagent. The iodoform test for methyl ketones uses ana queous alkaline iodine solution.

Spectroscopy

Tens of thousands of strong spectral lines in the 500–700 nm wavelength region make up the iodine molecule’s (I₂) spectrum. As a result, it is a frequently used secondary standard (wavelength reference).

The hyperfine structure of the iodine molecule is revealed by concentrating on one of these lines and detecting using a spectroscopic Doppler-free method.

We now resolve a line such that its components are either measurable (15 from even rotational quantum numbers, Jeven) or measurable (21 from odd rotational quantum numbers, Jodd).

Crystal scintillators are used to detect gamma rays using sodium iodide doped with thallium and calcium iodide. The resolution is not very good, but energy dispersive spectroscopy can be done with a high efficiency.

Spacecraft Propulsion

When compared to gridded ion thrusters used to propel earlier spacecraft, such as NASA’s DART mission, ESA’s GOCE satellite, or Japan’s Hayabusa probes, propulsion systems using iodine as the propellant can be built more compactly, with less mass (and cost), and operate more efficiently.

These systems all used xenon as the reaction mass. Iodine ions are a viable replacement for xenon because of their atomic weight, which is only 3.3% less than that of xenon, and their first two ionization energies, which are on average 12% less.

Applying iodine should enable ion-thrust technology to be used more widely, especially for smaller-scale spacecraft.

The European Space Agency says, “This small but potentially disruptive innovation could help to clear the skies of space junk.

by making it possible for tiny satellites to quickly and cheaply self-destruct at the conclusion of their missions by guiding themselves into the atmosphere and burning up there.

“The French organization ThrustMe carried out an in-orbit test of an electric-powered ion thruster for spacecraft at the beginning of 2021.

Iodine was utilized as the plasma source instead of xenon in order to accelerate ions with an electrostatic field and produce thrust.

Medicine

Elemental Iodine

When iodide is added to elemental iodine that is poorly soluble in water, an in situ reaction occurs that releases some free elemental iodine that can be used for antisepsis. Either way, elemental iodine is employed as an antiseptic.

Iodine shortage can also be treated with elemental iodine. An alternate method of producing iodine is to use iodophors, which are solutions of iodine complexed with a solubilizing agent (the iodide ion can be roughly interpreted as the iodophor in triiodide water solutions). Examples of these arrangements are as follows:

• Iodine tincture: either pure iodine in ethanol or iodine combined with sodium iodide in an ethanol and water mixture.
• Lugol’s iodine is primarily triiodide when iodine and iodide are combined in water alone. In contrast to tinctures of iodine, Lugol’s iodine contains less free iodine (I₂).
• Iodophor Povidone Iodine.
• Iodine-V: Iodine (I₂) and fulvic acid combine to form a clathrate compound, in which fulvic acid acts as a “cage” for iodine molecules. a crystalline,

solid, stable, and water-soluble compound. Iodine-V exclusively contains iodine in molecular (I2) form, in contrast to other iodophors. Operating rooms utilize iodine because of its fast antibacterial effect and low concentration of the drug. Its precise mechanism of action is unclear.

When it enters bacteria, it targets specific amino acids (such as cysteine and methionine), nucleotides, and fatty acids, which eventually cause cell death.

Although parvoviruses and nonlipid viruses are less vulnerable to it than lipid-enveloped viruses, it nevertheless possess antiviral properties.

Iodine most likely targets the surface proteins of viruses that are enclosed, and it also has the potential to react with unsaturated carbon bonds to destabilize membrane fatty acids.

Other Formulations

Iodide salts were administered orally in the past to treat lead or mercury poisoning before the development of organic chelating agents. Loui Melsens and other physicians from the 19th and early 20th centuries greatly popularized this method.

In medicine, acute thyrotoxicosis is treated with a saturated potassium iodide solution. When iodine-131 is utilized in radiopharmaceuticals (like iobenguane) that are not intended for the thyroid or tissues similar to the thyroid, it is also used to prevent uptake of the isotope in the thyroid gland (refer to the isotopes section above).

Nuclear fallout contains iodine-131, which is typically in the form of iodide. This is especially harmful because ingested iodine is concentrated by the thyroid gland and retained for longer lengths of time than the radiological half-life of eight days for this isotope.

This is why non-radioactive potassium iodide tablets may be recommended to those who are at risk of exposure to ambient radioactive iodine (iodine-131) in fallout.

One 130 mg tablet per 24 hours, or 100 mg (100,000 micrograms) of ionic iodine, is the usual adult dosage. (For normal health, the recommended daily intake of iodine is around 100 micrograms; refer to below).

This high dose of non-radioactive iodine reduces the thyroid gland’s ability to absorb radioactive iodine.

The radioactive half-life of iodine-131, which is typically present as iodide, makes it a particularly hazardous component of nuclear fallout because of the thyroid gland’s tendency to concentrate ingested iodine and hold onto it for extended periods of time. People who could be exposed to radioactive iodine (iodine-131) in the environment as a result of fallout may be advised to take potassium iodide pills that are not radioactive.

One 130mg tablet, or 100,000 micrograms, of ionic iodine, is the standard adult dosage, to be taken once every 24 hours.

(For optimal health, an individual should consume around 100 micrograms of iodine per day. below.) The thyroid gland absorbs less radioactive iodine when this high dosage of non-radioactive iodine is ingested.

Iodine is an element with a high atomic number and electron density. Because of the photoelectric effect of its innermost electrons, it may absorb X-rays with energies lower than 33.3 keV.

Intravenous injection of organoiodine compounds is utilized as an X-ray radiocontrast agent. This technology is frequently used in tandem with sophisticated X-ray methods like CT scanning and angiography. Currently, all radiocontrast materials that dissolve in water depend on substances that contain iodine.

Others

The majority of the iodine that is available is used in the synthesis of ethylenediamine dihydroiodide, which is fed to cattle as a dietary supplement.

The catalyst for the synthesis of acetic acid by the Monsanto and Cativa processes is another important application. The methanol feedstock in these technologies is converted by hydroiodic acid into methyl iodide, which then goes through carbonylation, meeting the global need for acetic acid.

Acetic acid is produced when the resultant acetyl iodide is hydrolyzed, regenerating hydroiodic acid. Certain applications exist for inorganic iodides.

The van Arkel–de Boer procedure, which includes the reversible synthesis of these elements’ tetraiodides, purifies thorium, zirconium, hafnium, and titanium. One of the main components of conventional photographic film is silver iodide.

Every year, thousands of kilograms of silver iodide are used to seed clouds in order to create rain. Erythrosine, an organoiodine molecule, is a significant food coloring ingredient.

Prototypes of significant surfactants, such as perfluorooctanesulfonic acid, are perfluoroalkyl iodides.

An example of a seemingly oscillating reaction (the concentration of an intermediate product is the only thing that oscillates) is the iodine clock reaction, a famous educational demonstration experiment in which iodine also functions as a starch test by creating a dark blue complex.

Despite the fact that iodine is widely distributed among many species, in an agricultural system, various species may be affected differently by agents containing iodine.

At quantities that do not affect the crop, an iodine-containing fungistatic (AJ1629-34EC) strongly inhibits the growth of all strains of Fusarium verticillioides.

Because of its comparatively natural chemistry, this anti-fungal agriculture therapy may be less toxic. In order to determine which ligands bind to which plant pattern recognition receptors (PRRs), ¹²⁵I is employed as the radiolabel.

Biological Role

As the heaviest element often required by living things, iodine (atomic number Z = 53) is a necessary component of life. A small number of microbes use tungsten (Z = 74), uranium (Z = 92), lanthanum, and other lanthanides.

The growth-regulating thyroid hormones thyroxine and triiodothyronine (T₄ and T₃, respectively, named for their number of iodine atoms) cannot be synthesized without it.

Iodine insufficiency causes the thyroid to produce less T₃ and T₄, which in turn causes the thyroid to expand in an attempt to get more iodine and ultimately results in the condition. referred to as a basic goiter.

Thyroxine (T₄), the predominant type of thyroid hormone in blood, has a longer half-life than T₃. The ratio of T₄ to T₃ that is released into the blood in humans ranges from 14:1 to 20:1.

T₄ is changed by deiodinases (5′-iodine) into active T₃, which is three to four times more effective than T₄. These are subsequently processed to yield iodothyronamine (T₁a) and thyroxine (T₀a’) by decarboxylation and deiodination.

Since the three deiodinase isoforms are selenium-containing enzymes, selenium from food is necessary for the synthesis of T₃.

Of the molecular weight of T₄ and T₃, iodine makes up 65% and 59%, respectively. Thyroid tissue and hormones contain 15–20 mg of iodine, whereas other tissues contain 70% of the total iodine in the body, including the cervix, eyes, stomach mucosa, fetal thymus, cerebrospinal fluid, artery walls, and salivary glands.

The placenta may store and release nutrients during pregnancy. assemble iodine. Iodide enters those tissues’ cells directly through the sodium-iodide symporter (NIS).

Iodine’s effects on fetal and newborn development are associated with mammary tissue; however, its effects on other tissues are (at least) little understood.

Dietary Recommendations and Intake

According to the US National Academy of Medicine, daily consumption levels should be between 110 and 130 µg for infants up to 12 months, 90 µg for kids up to eight years old, 130 µg for kids up to 13 years old, 150 µg for adults, 220 µg for expectant mothers, and 290 µg for nursing. Adults can tolerate up to 1,100 μg of exposure each day.

By examining the impact of supplementation on thyroid-stimulating hormone, this upper limit was determined.

The combined collection of data is referred to as Dietary Reference Values by the European Food Safety Authority (EFSA); AI and UL are defined in the same way as in the US, and Population Reference Intake (PRI) is used in place of RDA and Average Requirement (EAR).

The PRI for iodine is 150 μg/day for men and women over the age of 18, and 200 μg/day for women who are pregnant or nursing. For kids ages 1 to 17, the PRI rises from 90 to 130 μg/day as they get older.

Except for the breastfeeding requirement, these PRIs are equivalent to the RDAs in the United States. The thyroid gland can synthesize the necessary daily levels of T₄ and T₃ with as little as 70 μg per day.

Higher recommended daily allowance levels of iodine appear to be required for several physiological systems, such as the choroid plexus, brain cells, salivary glands, thymus, breastfeeding, and artery walls, to function at their best.

As long as the animals get enough iodine, dairy products, eggs, fish, shellfish, seaweeds (like kelp), and plants produced in iodine-rich soil are natural sources of dietary iodine.

Iodine, in the form of potassium or sodium iodate, is added to iodized salt to fortify it. In the United States, the median daily consumption of iodine from food was 190–210 μg for women and 240–300 μg for men as of 2000.

While pregnant and childbearing women may be at minor risk for iodine shortage, the overall US population has an adequate iodine diet.

The amount consumed in Japan was thought to be significantly larger, ranging from 5,280 μg to 13,800 μg per day from dietary seaweed or kombu kelp, frequently in the form of kombu umami extracts for potato chips and soup stock.

New research, however, points to Japan’s consumption being closer to 1,000–3,000 μg/day. In Japan, the adult UL was last updated in 2015 and is now 3,000 µg/day.

There have been occasional reports of iodine-induced hyperthyroidism following the implementation of iodine fortification initiatives, such as the iodization of salt (also known as the Jod-Basedow phenomenon).

The illness tends to primarily affect individuals over forty, and acute iodine deficiency and a large initial spike in iodine intake seem to be associated with increased risk.

Deficiency

Extreme fatigue, goiter, mental slowing, depression, weight gain, and low basal body temperatures are symptoms of hypothyroidism, which is caused by low dietary iodine levels in semi-arid equatorial climates and remote inland areas where no marine foods are consumed.

The most common preventable cause of intellectual disability in children is iodine deficiency, which generally manifests as hypothyroidism in infants or young children.

In wealthy countries, this issue has largely been resolved by adding iodine to table salt. Iodine deficiency, however, continues to be a significant public health issue in underdeveloped countries.

In various parts of Europe, iodine shortage is also an issue. When children with moderate iodine deficiency receive a replacement, their visual problem-solving, fine motor skills, and information processing all improve.

Precautions

Toxicity

If elemental iodine (I₂) is consumed undiluted, it can be hazardous. A human weighing 70–80 kg would require a deadly dose of 30 mg/kg, or roughly 2.1–2.4 kilos, even though studies on rats showed that these animals could survive after consuming a dose of 14000 mg/kg.

When selenium deficiency is present, excess iodine may be more harmful. This is one of the theoretical reasons why iodine supplementation in people weak in selenium could be troublesome.

Its oxidizing characteristics, which cause it to denaturize proteins (including enzymes), are the source of its toxicity. Iodine elements can also irritate the skin.

Solid iodine crystals should be handled carefully since they can cause harm if they come into direct contact with the skin.

High elemental iodine concentration solutions, like tincture of iodine and Lugol’s solution, can damage tissue when used for extended cleaning or antisepsis; in certain documented cases, liquid Povidone-iodine (Betadine) trapped against the skin caused chemical burns.

Occupational Exposure

Iodine exposure in the workplace can occur through skin contact, ocular contact, ingestion, and inhalation.

The legal limit (permissible exposure limit) for iodine exposure at work is set by the Occupational Safety and Health Administration (OSHA) at 0.1 ppm (1 mg/m³) for an 8-hour workday.

A Recommended Exposure Limit (REL) of 0.1 ppm (1 mg/m³) has been established by the National Institute for Occupational Safety and Health (NIOSH) for an 8-hour workday. Iodine poses an urgent risk to life and health at values of 2 ppm.

Allergic Reactions

Foods and items containing iodine might cause hypersensitivity reactions in certain persons. Iodine or betadine tinctures can sometimes result in severe rashes when applied topically. Anaphylaxis can be fatal when iodine-based contrast agents are used parenterally (see above).

Reactions can range from a modest rash to this. Due to these symptoms, there is a common misconception—even among doctors—that some patients have an allergy to iodine; sensitivities to seafood that is high in iodine have also been misinterpreted.

Actually, no one has ever reported a case of a real iodine allergy; instead, allergies to elemental iodine or it is theoretically impossible to have simple iodide salts.

A person who has shown an allergy to one food or product containing iodine may not have an allergic reaction to another.

Hypersensitivity reactions to meals and goods containing iodine appear to be related to their other molecular components. Individuals who are allergic to any number of foods (shellfish, eggs, milk, etc.) are not more likely to experience hypersensitivity to contrast media.

Before administering any treatment containing iodine, like with all medications, it is important to inquire about and review the patient’s allergy history.

US DEA List I status

The conversion of elemental iodine to hydroiodic acid by phosphorus is an efficient method of reducing ephedrine or pseudoephedrine to methamphetamine.

Because of this, iodine was listed as a List I precursor chemical under 21 CFR 1310.02 by the US Drug Enforcement Administration.

FAQ

Author: Arjun Sharma