Halide
In chemistry, a halide (rarely halogenide[1]) is a binary chemical compound consisting of a halogen atom and an element or radical that is less electronegative (or more electropositive) than the halogen, to make a fluoride, chloride, bromide, iodide, astatide, or theoretically tennesside compound. The alkali metals combine directly with halogens under appropriate conditions, forming halides of the general formula MX (X = F, Cl, Br or I). Many salts are halides; the hal- syllable in halide and halite reflects this correlation.
A halide ion is a halogen atom bearing a negative charge. Naturally occurring halide anions include fluoride (F−), chloride (Cl−), bromide (Br−), and iodide (I−), while astatide (At−) can be produced by reducing free astatine.[2] Halide ions are present in all ionic halide salts and halide minerals.
Halides are colourless, high melting crystalline solids with highly negative enthalpies of formation.
History
[edit]The name "halogen" - derived from Greek roots for "sea salt" (hal-) and "to produce" (-gen), indicates the ability of these elements to form binary compounds such as sea salt with similar properties. These binary compounds, normally consisting of a halide anion and metal cation, are the most common occurance of halides.[3]
Halide compounds have been well-documented for much of human history; written records as old as the Bible contain reference to the importance of rock salt (NaCl) with evidence of its use extending much further, and medieval scientists used mixtures of hydrochloric acid (HCl) with nitric acid to dissolve gold.[3]
Due to their reactivity, halogen species naturally occur almost exclusively as halide anions. Because of this, common halide compounds were used to isolate several halogen species for the first time. In 1774, Carl Wilhelm Scheele reacted sodium chloride (NaCl) with sulfuric acid and manganese dioxide to isolate and identify elemental chlorine gas for the first time. Similarly, Henri Moissan prepared elemental fluorine in 1886 by the electrolysis of potassium hydrogen fluoride (KHF2) in hydrofluoric acid (HF)[4].
Properties
[edit]Generally, there are four types of Halides:
- Hydrogen Halides: (Hydrohalic acids in aqueous phase)
Halides with the general formula HX (e.g. HF, HCl, HBr, HI). All known hydrogen halides are gases at room temperature; however, they appear more commonly in the aqueous phase as hydrohalic acids. These binary acids are strong because their halogen ions are very soluble in water.
- Alkali & Alkaline Earth Metal Halides:
Halides with the general Formula LiXj, where L is an element from group 1 or 2 on the periodic table (e.g. LiCl, NaCl, NaBr, MgBr2, CaCl2). These compounds are ionic salts and are soluble in water. They are solid at room temperature and have very high melting points.[5]
Halides with the general formula MiXj, where M is an element from group 3 - 11 on the periodic table (e.g. AgCl, CuCl, CuCl3, TiCl4). These compounds are usually soluble in water; however, since the compound is coordinated with a metal, water ligands can bind to the complex to form hydrates. They usually have high melting points. Metals also have different oxidation states, meaning a matrix of different i & j values can be produced[6]. For example, CrCl2, CrCl3, CrF4, and CrF5 are all valid compounds with possible different coordination geometries.
- Non-Metal Halides:
Halides with the general Formula NiXj, where N is an element from group 12 - 18 on the periodic table (e.g. AlCl3, SiCl4, OF2, PCl3, PCl5, XeF4). Usually, these compounds are not soluble in water. Some non-metal halides will react with water due to their relative instability. They usually have relatively lower melting points. The properties of non-metals vary, as non-metal elements have a wide range of properties.[6]
Halide Structure Arrangements
[edit]In solid phase, halides can arrange themselves into the following intermolecular geometries:
- Crystalline Halides
Some halides form lattice structures, in which ionic forces lock molecules in place, orienting them towards their opposing charge. Lattices can take many forms; NaCl takes a cubic lattice form.[8]
In liquid and gas phase, crystalline halides often decompose into discrete halide structures. Some structures stabilize by bonding to multiples of these discrete halide structures known as molecular halides. For example, AlCl3 is a molecular halide because its aluminum bonds with another chlorine atom from another monomer and vice versa to form a tetramolecular bond. It therefore takes form as Al2Cl6 in room temperature.[6]
- Polymeric (Chain-like) Halides
Similar to molecular halides, structures may stabalize by bonding to other monomers. Some compounds continue bonding to produce chains of arbitrary length, and can span 1, 2, or 3 dimensions. For example, BeCl2 at high temperatures breaks into monomers, but at room temperature solidifies into a 1-dimensional polymeric halide. The chlorines bridge the compounds together so that electrons can flow more freely, which further stabalizes the system.[6]
- Discrete Halide Structures
Few halides are unable to produce lattice or polymeric structures for varying reasons, such as UF6.[9]
Pseudohalides are similar to halides, however, they contain a pseudohalogen - an anionic compound that is quite structurally stable such that it has near-identical properties to a halogen. For example, silver cyanide (AgCN) can react similarly to silver halides because the cyanide ion CN- is a pseudohalogen.[6] Pseudohalogen properties are not standardized, meaning not every pseudohalogen can be used in place of a halogen for every reaction and vice versa.
Tests
[edit]Silver Nitrate Test
[edit]_colour_test_with_AgNO3.jpg)
Dilute nitric acid can be added to acidify a solution to obtain a solution of silver nitrate and dilute nitric acid. Nitric acid removes any other ions that may form precipitates with silver nitrate[10]. This solution is then added to the compound being tested to yield the following results:
| ion present | observation |
|---|---|
| F- | no precipitate |
| Cl- | white precipitate |
| Br- | cream (pale) precipitate |
| I- | yellow (pale) precipitate |
The precipitates formed in this reaction are insoluble silver halides:
Ag+aq + Cl−(aq) → AgCl(s)
Ag+aq + Br−(aq) → AgBr(s)
Ag+aq + I−(aq) → AgI(s)
Silver fluoride is soluble, explaining why no precipitate forms:
Ag+aq + F−(aq) → Ag+aq + I−(aq)
Additional Ammonia Test
[edit]After producing precipitates using nitric acid, both dilute and concentrated ammonia solution can be used to confirm the ions present. This yields the following results:
| Precipitate | With Dilute Ammonia Solution | With Concentrated Ammonia Solution |
|---|---|---|
| AgCl | Precipitate dissolves (colourless solution formed) | Precipitate dissolves (colourless solution formed) |
| AgBr | Precipitate remains (no change) | Precipitate dissolves (colourless solution formed) |
| AgI | Precipitate remains (no change) | Precipitate remains (no change) |
For silver halides, the solubility product is expressed as:
Ksp = [Ag+][X−]
Where the square brackets are used to represent molar concentrations (mol L-1).
These values can be compared against the solubility products of the compounds shown below:
| Compound | Ksp (mol2dm-6) |
|---|---|
| AgCl | 1.8 x 10-10 |
| AgBr | 7.7 x 10-13 |
| AgI | 8.3 x 10-17 |
If a calculated Ksp value is less than the solubility product, no precipitate is formed.[11] If a calculated Ksp value is greater than the solubility product, a precipitate is formed. The solubility product of AgF cannot be reported because it is too soluble.
Concentrated Sulfuric Acid Test
[edit]Concentrated sulfuric acids can be added to solids to test for halides. This test yields the following results:
| Ion Present | Observation |
|---|---|
| F- | steamy acidic fumes (HF) |
| Cl- | steamy acidic fumes (HCl) |
| Br- | steamy acidic fumes (HBr) with brown bromine vapour |
| I- | some HI fumes with purple iodine vapour (+ a red compound in reaction vessel) |
This test cannot be used to distinguish fluoride and chloride, so a silver nitrate test would be needed instead.
Uses
[edit]Metal halides are used in high-intensity discharge lamps called metal halide lamps, such as those used in modern street lights. These are more energy-efficient than mercury-vapor lamps, and have much better colour rendition than orange high-pressure sodium lamps. Metal halide lamps are also commonly used in greenhouses or in rainy climates to supplement natural sunlight.
Silver halides are used in photographic films and papers. When the film is developed, the silver halides which have been exposed to light are reduced to metallic silver, forming an image.
Sulfur halides, specifically sulfur chlorides, are used in rubber for vulcanization, and are precursors for the synthesis of mustard gas.[12]
Halides are also used in solder paste, commonly as a Cl or Br equivalent.[13]
Synthetic organic chemistry often incorporates halogens into organohalide compounds.
See also
[edit]References
[edit]- ^ "Definition of HALOGENIDE". www.merriam-webster.com. Retrieved 2022-01-07.
- ^ Busch, Marianna Anderson (2003-01-01), Meyers, Robert A. (ed.), "Halogen Chemistry", Encyclopedia of Physical Science and Technology (Third Edition), New York: Academic Press, pp. 197–222, ISBN 978-0-12-227410-7, retrieved 2024-11-05
- ^ a b Busch, Marianna Anderson (2003-01-01), Meyers, Robert A. (ed.), "Halogen Chemistry", Encyclopedia of Physical Science and Technology (Third Edition), New York: Academic Press, pp. 197–222, ISBN 978-0-12-227410-7, retrieved 2024-11-05
- ^ Emsley, John (2011). Nature's building blocks: everything you need to know about the elements (New ed., completely rev. and updated ed.). Oxford: Oxford University Press. ISBN 978-0-19-960563-7.
- ^ Huheey, James E.; Keiter, Ellen A.; Keiter, Richard L. (1993). Inorganic chemistry: principles of structure and reactivity (4th ed.). New York, NY: HarperCollins College Publishers. ISBN 978-0-06-042995-9.
- ^ a b c d e f "4.6: Halogens and Halides". Chemistry LibreTexts. 2018-12-26. Retrieved 2024-11-05.
- ^ Tramsek, Melita; Zemva, Boris (2007-05-22). "Synthesis, Properties and Chemistry of Xenon(II) Fluoride". ChemInform. 38 (21). doi:10.1002/chin.200721209. ISSN 0931-7597.
- ^ Wiley-VCH, ed. (2003-03-11). Ullmann's Encyclopedia of Industrial Chemistry (1 ed.). Wiley. doi:10.1002/14356007.a24_317.pub4. ISBN 978-3-527-30385-4.
- ^ a b Scott, Robert A., ed. (2012-02-01). Encyclopedia of Inorganic and Bioinorganic Chemistry (2 ed.). Wiley. doi:10.1002/9781119951438.eibc0078.pub2. ISBN 978-1-119-95143-8.
- ^ "Testing for halide ions". www.chemguide.co.uk. Retrieved 2024-11-05.
- ^ "Testing for Halide Ions". Chemistry LibreTexts. 2013-10-03. Retrieved 2024-11-05.
- ^ Dirican, Dilcan; Pfister, Nils; Wozniak, Martin; Braun, Thomas (2020-06-02). "Reactivity of Binary and Ternary Sulfur Halides towards Transition-Metal Compounds". Chemistry – A European Journal. 26 (31): 6945–6963. doi:10.1002/chem.201904493. ISSN 0947-6539. PMC 7318666. PMID 31840851.
- ^ "Halogen-Free Solder Paste" (PDF). Archived from the original (PDF) on 2012-03-17. Retrieved 2011-03-21.
