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2,4,6-Tribromoaniline is a brominated derivative of aniline with the formula C 6 H 4 Br 3 N. It is used in organic synthesis of pharmaceuticals, agrochemicals and fire-extinguishing agents. It is used in organic synthesis of pharmaceuticals, agrochemicals and fire-extinguishing agents.
Brominating aniline with elemental bromine gives 2,4,6-tribromoaniline. This is then diazotized, then reacted with ethanol to replace the diazonium group with hydrogen, forming 1,3,5-tribromobenzene. [3] It has also been prepared by these methods: [3] replacement of the amino group of 3,5-dibromoaniline with bromine
Anthropogenic and natural sources of bromine. The major sources include sea spray, salt lakes, marshes, volcanos, anthropogenic sources. Sinks include exchange of brominated compounds with the stratospheric and troposphere.Bromine's chemistry is linked to other halogens such as chlorine and iodine amplify atmospheric cycling that contributes to troposphere and stratosphere ozone layer ...
The bromoanilines form a group of three isomers where the bromine atom occupies the para, ortho or meta position on the aromatic ring. Bromoaniline isomers Arene substitution patterns. The three isomers are: 2-Bromoaniline (o-Bromoaniline) [1] 3-Bromoaniline (m-Bromoaniline) [2] 4-Bromoaniline (p-Bromoaniline) [3]
Bromine is more electronegative than carbon (2.9 vs 2.5). Consequently, the carbon in a carbon–bromine bond is electrophilic, i.e. alkyl bromides are alkylating agents. [2] Carbon–halogen bond strengths, or bond dissociation energies are of 115, 83.7, 72.1, and 57.6 kcal/mol for bonded to fluorine, chlorine, bromine, or iodine, respectively ...
4-Bromoaniline is a compound where an aniline molecule is substituted with a bromine atom on the para position. Commercially available, this compound may be used as a building block, e.g. in the preparation of monobrominated biphenyl via the Gomberg-Bachmann reaction. [2]
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Bond energies to bromine tend to be lower than those to chlorine but higher than those to iodine, and bromine is a weaker oxidising agent than chlorine but a stronger one than iodine. This can be seen from the standard electrode potentials of the X 2 /X − couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, approximately +0.3 V ...