# General Chemistry/Chemistries of Various Elements/Group 17

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## Halogens

The halogens are found in Group 17(VIIA). The halogens are fluorine, chlorine, bromine, iodine, and astatine.
Like the alkali metals, the halogens are extremely reactive. They have seven valence electrons, meaning they require only one more electron for a noble configuration. This gives them very large electron affinities and extreme reactivity to form ions with a -1 charge. They are so reactive that in their homogeneous state, UV light will catalyze a radical reaction.
The halogens exist in diatomic form. Under normal conditions, they will always occur in pairs, covalently bonded. The covalent bond allows them to share an electron and possess a complete octet. F$_2$ is a pale yellowish-brown gas. It is highly reactive, causing organic compounds and hydrogen gas to explode, even without a spark. Cl$_2$ is a pale yellow-green gas. It reacts with water to form disinfectants and bleaches. Br$_2$ is a reddish-brown liquid, but, being volatile, it readily evaporates into a reddish vapor. I$_2$ is a gray solid that forms a violet gas if heated.
Fluorine is the most electronegative of all elements, and it is so reactive that it attacks almost any other element (noble gases, oxygen, nitrogen, and gold are the exceptions) to form fluorides. Chlorine is somewhat less reactive, bromine somewhat less reactive than chlorine, and iodine even less, but even iodine is a formidable ionizer. Extreme radioactivity masks the chemical properties of astatine. With increasing atomic weight for these elements, the elements have higher boiling and melting points. At normal temperatures, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid.
* The Halogens
* Copper (II) Fluoride
* Chlorine gas
* Bottle containing both liquid and gaseous bromine
* Solid iodine pieces

## Reactions

| $2{\hbox{F}}_{2(g)}+2{\hbox{H}}_{2}{\hbox{O}}_{(l)}\to {\hbox{O}}_{2(g)}+4{\hbox{HF}}_{(aq)}$ ${\hbox{Cl}}_{2(g)}+{\hbox{H}}_{2}{\hbox{O}}_{(l)}\to {\hbox{HCl}}_{(aq)}+{\hbox{HClO}}_{(aq)}$ ${\hbox{Br}}_{2(g)}+{\hbox{H}}_{2}{\hbox{O}}_{(l)}\to {\hbox{HBr}}_{(aq)}+{\hbox{HBrO}}_{(aq)}$ | The halogens in diatomic form react with water to produce acids. Iodine does not react with water and is only slightly soluble. Chlorine also has low solubility, but it will react in water to form hypochlorous acid and hydrochloric acid. |
| ${\hbox{X}}_{2}+{\hbox{H}}_{2(g)}\to 2{\hbox{HX}}_{(g)}$ | In chemical reactions, an X is used to symbolize any halogen element. All halogens form gaseous compounds with hydrogen: hydrogen fluoride HF, hydrogen chloride HCl, hydrogen bromide HBr, and hydrogen iodide HI. These are acidic, strongly reactive substances called hydrogen halides. |
| ${\hbox{HX}}_{(g)}+2{\hbox{H}}_{2}{\hbox{O}}_{(g)}\to {\hbox{H}}_{3}{\hbox{O}}_{(aq)}^{+}+{\hbox{X}}_{(aq)}^{-}+{\hbox{H}}_{2}{\hbox{O}}_{(l)}$ | When hydrogen halides are dissolved in water, they are known as hydrohalic acids. Except for hydrogen fluoride, they are among the strongest known acids. |
| $2{\hbox{M}}_{(s)}+{\hbox{X}}_{2(g)}\to 2{\hbox{MX}}_{(s)}$ ${\hbox{HX}}_{(g)}+{\hbox{NH}}_{3(g)}\to {\hbox{NH}}_{4}{\hbox{X}}_{(g)}$ | These reactions show reactions with metals and ammonia gas to form salts and ammonium halides, respectively. |
| ${\hbox{SiCl}}_{4}+2{\hbox{H}}_{2}{\hbox{O}}\to {\hbox{SiO}}_{2}+4{\hbox{HCl}}$ | Many non-metallic halides react with water to release hydrogen halides. For example, silicon chloride and water react. |

## Other Compounds

Halides of metals are known as salts. Sodium chloride, better known as "table salt", is the crystalline substance often used to enhance the flavor of food. Note, however, that not all salts are halides (for example, sodium sulfate Na$_2$SO$_4$), and not all halides are salts (carbon tetrachloride, CCl$_4$).
Interhalogens are molecules composed of two or more different halogen atoms. They are similar to the diatomic halogens. Some examples are chlorine monofluoride ClF and bromine monochloride BrCl. There are many others, and they are all very reactive and somewhat unstable. Interhalogens take the form XY$_n$, where n is 1, 3, 5, or 7. X and Y are both halogens, X being the less electronegative. Larger molecules and ions are also known. E.g. I$_2$ does not dissolve in water very well, but it dissolves readily in a solution containing iodide ions under the formation of the triiode anion I$_2$ + I$^-$ --> I$_3$$^-$
Noble gas compounds have been formed using fluorine. Although noble gases are supposedly inert, the larger ones like xenon will form covalent bonds with a very electronegative element like fluorine. Xenon difluoride XeF$_2$, xenon tetrafluoride XeF$_4$, and xenon hexafluoride XeF$_6$ are among the noble gas compounds that have been created.

## Oxidation states

Fluorine is by far the most electronegative element and when it reacts it invariably assumes an oxidation number of -1. The compounds formed this way are called fluorides.
The other halogens ofter react in similar fashion, but they are less electronegative than fluorine. If they react in the presence of an element with a higher electronegativity than their own other (higher) oxidation numbers may result.
For chlorine this can only happen with two elements: oxygen and fluorine. The best known compounds with higher oxidation states are the oxo-acids and their salts.
| Ox # | Acid | Name | Anion | Name |
| +7 | HClO$_4$ | perchloric acid | ClO$_4$$^-$ | perchlorate ion |
| +5 | HClO$_3$ | chloric acid | ClO$_3$$^-$ | chlorate ion |
| +3 | HClO$_2$ | chlorous acid | ClO$_2$$^-$ | chlorite ion |
| +1 | HClO | hypochlorous acid | ClO$^-$ | hypochlorite ion |
| -1 | HCl | hydrochloric acid | Cl$^-$ | chloride ion |
As shown the oxidation numbers can range over a full octet from -1 to +7, representing an electron configuration for the chlorine atom ranging from [Ar] to [Ne]. The naming is somewhat makeshift. The endings -ic for the acid and -ate for the anion are typically reserved for the highest oxidation number, but they had already been given to the +5 species when the +7 ones were found. This is why the per- suffix was added. The +3 species have the usual -ous and -ite endings for a oxidation state one step below the highest and the hypo- prefix indicates that it is a an even lower oxidation state.
Some of these compounds are rather common. Hydrochloric acid can be found in our stomachs and as a household cleaner. In pure form it is a gas, but it is quite soluble in water. Table salt is a chloride salt, the hypochlorite ion is the active ingredient in bleach. Perchlorates are very soluble salts and this fact is often exploited if we wish to make a certain cation readily soluble.
The elements bromine and iodine form similar acids and salts to chlorine.

## Halo-acids

Oxo-acids like nitric acid HNO$_3$ of sulfuric acid H$_2$SO$_4$ and their anions like the nitrate ion NO$_3$$^-$ or the sulfate ion SO$_4$$^{2-}$ are well known. In these molecular ions oxygen functions as the more electronegative ligand around a central atom (N or S) in a high oxidation state.
Halogens being rather electronegative themselves can fulfill the same role as oxygen and this gives rise to halo-acids and their anions. A good example of such an anion is the tetrafluoroborate anion BF$_4$$^-$.
The naming of such compounds is derived from the oxo ones. The name borate with its -ate ending indicates that boron in its highest oxidation state is the central atom. The fact that the ligands are F rather than oxygen is indicated by the prefix fluoro- and the tetra- indicates that there are four ligands around the central atom. Similar names pertain to the other halogens, e.g. ZnCl$_4$$^{2-}$ is the tetrachlorozincate ion.