CAS 7553-56-2 Iodine Ball

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Iodine ball cas 7553-56-2

Iodine ball CAS 7553-56-2

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Iodine Ball CAS 7553-56-2 Description:

• Product Name: Iodine
• Cas No: 7553-56-2
• EINECS No: 231-442-4
• Molecular Formula: I2
• Molecular Weight: 253.809
• Melting Point: 113 °C (lit.)
• Boiling Point: 184 °C (lit.)
• Density: 1.32 g/mL at 25 °C
• Purity: 99%
• Appearance: purple black spherical
• Package: according to the client’s requirement
• production Capacity: 1-1000 Metric Ton/Day

Charge-transfer complexes:

The iodine molecule, I₂, dissolves in CCl₄ and aliphatic hydrocarbons to give bright violet solutions. In these solvents, the absorption band maximum occurs in the 520 – 540 nm region and is assigned to a π^* to σ^* transition. When I₂ reacts with Lewis bases in these solvents a blue shift in I₂ peak is seen and a new peak (230 – 330 nm) arises that is due to the formation of adducts, which are referred to as charge-transfer complexes.

Hydrogen iodide:

The simplest compound of iodine is hydrogen iodide, HI. It is a colourless gas that reacts with oxygen to give water and iodine. Although it is useful in iodination reactions in the laboratory, it does not have large-scale industrial uses, unlike the other hydrogen halides. Commercially, it is usually made by reacting iodine with hydrogen sulfide or hydrazine:

2 I₂ + N₂H₄ H₂O⟶ 4 HI + N₂

At room temperature, it is a colourless gas, like all of the hydrogen halides except hydrogen fluoride, since hydrogen cannot form strong hydrogen bonds to the large and only mildly electronegative iodine atom. It melts at −51.0 °C and boils at −35.1 °C. It is an endothermic compound that can exothermically dissociate at room temperature, although the process is very slow unless a catalyst is present: the reaction between hydrogen and iodine at room temperature to give hydrogen iodide does not proceed to completion. The H–I bond dissociation energy is likewise the smallest of the hydrogen halides, at 295 kJ/mol.

Aqueous hydrogen iodide is known as hydroiodic acid, which is a strong acid. Hydrogen iodide is exceptionally soluble in water: one litre of water will dissolve 425 litres of hydrogen iodide, and the saturated solution has only four water molecules per molecule of hydrogen iodide. Commercial so-called “concentrated” hydroiodic acid usually contains 48–57% HI by mass; the solution forms an azeotrope with boiling point of 126.7 °C at 56.7 g HI per 100 g solution. Hence hydroiodic acid cannot be concentrated past this point by evaporation of water.

Unlike hydrogen fluoride, anhydrous liquid hydrogen iodide is challenging to work with as a solvent, because its boiling point is low, it has a small liquid range, its dielectric constant is low and it does not dissociate appreciably into H₂I⁺ and HI⁻
₂ ions – the latter, in any case, are much less stable than the bifluoride ions (HF⁻
₂) due to the very weak hydrogen bonding between hydrogen and iodine, though its salts with very large and weakly polarising cations such as Cs⁺ and NR⁺
₄ (R = Me, Et, Buⁿ) may still be isolated. Anhydrous hydrogen iodide is a poor solvent, able to dissolve only small molecular compounds such as nitrosyl chloride and phenol, or salts with very low lattice energies such as tetraalkylammonium halides.

Other binary iodides:

Nearly all elements in the periodic table form binary iodides. The exceptions are decidedly in the minority and stem in each case from one of three causes: extreme inertness and reluctance to participate in chemical reactions (the noble gases); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyond bismuth); and having an electronegativity higher than iodine’s (oxygen, nitrogen, and the first three halogens), so that the resultant binary compounds are formally not iodides but rather oxides, nitrides, or halides of iodine. (Nonetheless, nitrogen triiodide is named as an iodide as it is analogous to the other nitrogen trihalides.)

Given the large size of the iodide anion and iodine’s weak oxidising power, high oxidation states are difficult to achieve in binary iodides, the maximum known being in the pentaiodides of niobium, tantalum, and protactinium. Iodides can be made by the reaction of an element or its oxide, hydroxide, or carbonate with hydroiodic acid, and then dehydrated by mildly high temperatures combined with either low-pressure or anhydrous hydrogen iodide gas.

These methods work best when the iodide product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative iodination of the element with iodine or hydrogen iodide, high-temperature iodination of a metal oxide or other halide by iodine, a volatile metal halide, carbon tetraiodide, or an organic iodide. For example, molybdenum(IV) oxide reacts with aluminium(III) iodide at 230 °C to give molybdenum(II) iodide. An example involving halogen exchange is given below, involving the reaction of tantalum(V) chloride with excess aluminium(III) iodide at 400 °C to give tantalum(V) iodide:

{\displaystyle {\ce {3TaCl5 + {\underset {(excess)}{5AlI3}}-> 3TaI5 + 5AlCl3}}}

Lower iodides may be produced either through thermal decomposition or disproportionation or by reducing the higher iodide with hydrogen or a metal, for example:

{\displaystyle {\ce {TaI5{}+Ta->[{\text{thermal gradient}}][{\ce {630^{\circ }C\ ->\ 575^{\circ }C}}]Ta6I14}}}

Most of the iodides of groups 1, 2, and 3, along with the lanthanides and actinides in the +2 and +3 oxidation states, are mostly ionic. At the same time, nonmetals tend to form covalent molecular iodides, as do metals in high oxidation states from +3 and above.

Ionic iodides MI_n tend to have the lowest melting and boiling points among the halides MX_n of the same element because the electrostatic forces of attraction between the cations and anions are weakest for the large iodide anion.

In contrast, covalent iodides tend to instead have the highest melting and boiling points among the halides of the same element, since iodine is the most polarisable of the halogens and, having the most electrons among them, can contribute the most to van der Waals forces.

Naturally, exceptions abound in intermediate iodides where one trend gives way to the other. Similarly, solubilities in water of predominantly ionic iodides (e.g. potassium and calcium) are the greatest among ionic halides of that element, while those of covalent iodides (e.g. silver) are the lowest of that element. In particular, silver iodide is very insoluble in water and its formation is often used as a qualitative test for iodine.