Halogens

Fig. 1 - Group 7 or group 17? Sometimes it is just easier to refer to them as 'the halogens'

• This article is an introduction to the halogens.
• We’ll look at their properties and characteristics before taking a closer look at each member in turn.
• We’ll then outline some of the reactions they take part in and their uses.
• Finally, we’ll also explore how you can test for the presence of halide ions in compounds.

Halogen properties

The halogens are all non-metals. They show many of the properties typical of non-metals.

• They are poor conductors of heat and electricity.
• They form acidic oxides.
• When solid, they are dull and brittle. They also sublime easily.
• They have low melting and boiling points.
• They have high electronegativity values. In fact, fluorine is the most electronegative element in the periodic table.
• They form anions, which are ions with negative charges. The first four halogens all commonly form anions with a charge of -1, meaning they have gained one electron.
• They also form diatomic molecules.

Fig. 2 -A diatomic chlorine molecule, made from two chlorine atoms

We call ions made from halogen atoms halides. Ionic compounds made from halide ions are called halide salts. For example, the salt sodium chloride is made from positive sodium ions and negative chloride ions.

Fig. 3 - A chlorine atom, left, and chloride ion, right

Trends in properties

Reactivity and electronegativity decrease going down the group whilst atomic radius and melting and boiling points increase. Oxidising ability decreases going down the group whilst reducing ability increases.

Halogens elements

At the start of this article, we said that the halogen group contains six elements. But it depends on who you ask. The first four members are known as the stable halogens. These are fluorine, chlorine, bromine, and iodine. The fifth member is astatine, an extremely radioactive element. The sixth is the artificial element tennessine, and you’ll find out why some people don’t include it in the group later on. Let’s now take a look at the elements individually, starting with fluorine.

Fluorine

Fluorine is the smallest and lightest member of the group. It has the atomic number 9, and is a pale yellow gas at room temperature.

Fluorine is the most electronegative element in the periodic table. This makes it one of the most reactive elements as well. This is because it is such a small atom. Halogens tend to react by gaining an electron to form a negative ion. Any incoming electrons feel a strong attraction to fluorine’s nucleus because the fluorine atom is so small. This means fluorine reacts readily. In fact, fluorine forms compounds with almost all other elements. It can even react with glass! We store it in special containers using metals such as copper, as they form a protective layer of fluoride on their surface.

Fluorine’s name comes from the Latin verb fluo-, meaning 'to flow', which reflects its origins. Fluorine was originally used to lower the melting points of metals for smelting. In the 1900s it was used in refrigerators in the form of CFCs, or chlorofluorocarbons, which are now banned due to their harmful effect on the ozone layer. Nowadays fluorine is added to toothpaste and is a part of Teflon™.

Fig-4 Liquid Fluorine in cryogenic bath, wikimedia commons[1]

Chlorine

Chlorine is the next smallest member of the halogens. It has an atomic number of 17 and is a green gas at room temperature. Its name comes from the Greek word chloros, meaning 'green'.

Chlorine has a pretty high electronegativity, behind only oxygen, and its close cousin fluorine. It is also extremely reactive and is never found naturally in its elemental state.

As we mentioned earlier, melting and boiling points increase as you move down the group in the periodic table. This means that chlorine has higher melting and boiling points than fluorine. However, it has a lower electronegativity, reactivity, and first ionisation energy.

We use chlorine for a wide range of purposes, from making plastics to disinfecting swimming pools. However, it is more than just a conveniently useful element. It is essential to life for all known species. But too much of a good thing can be bad, and this is exactly the case with chlorine. Chlorine gas is highly toxic, and was first used as a weapon in World War One.

Fig .5- An ampoule of chlorine gas, W.Oelen, Wikimedia commons [2]

Bromine

The next element is bromine. Bromine is a dark red liquid at room temperature, and has an atomic number of 35.

Like fluorine and chlorine, bromine does not occur freely in nature but instead forms other compounds. These include organobromides, which we commonly use as fire retardants. Over half of the bromine produced globally each year is used in this way. Like chlorine, bromine can be used as a disinfectant. However, chlorine is preferred due to bromine’s higher cost.

Fig. 6- An ampoule of liquid bromine, Jurii, CC BY 3.0, wikimedia commons [3]

Iodine

Iodine is the heaviest of the stable halogens, with an atomic number of 53. It is a grey-black solid at room temperature and melts to produce a violet liquid. Its name comes from the Greek iodes, meaning 'violet'.

The trends outlined earlier in the article continue as you move down the periodic table to iodine. For example, iodine has a higher boiling point than fluorine, chlorine, and bromine, but a lower electronegativity, reactivity, and first ionisation energy. However, it is a better reducing agent.

Fig. 7 - A sample of solid iodine. commons.wikimedia.org, Public domain

Astatine

Now we come to astatine. This is where things start to get a little more interesting.

Astatine has an atomic number of 85. It is the rarest naturally occurring element in the Earth’s crust, mostly found leftover as other elements decay. It is pretty radioactive - its most stable isotope only has a half-life of just over eight hours!

A sample of pure astatine has never been successfully isolated because it would immediately vaporise under the heat of its own radioactivity. Because of this, scientists have had to make guesses about most of its properties. They predict it follows the trends shown in the rest of the group, and so give it a lower electronegativity and reactivity than iodine, but higher melting and boiling points. However, astatine also shows some unique properties. It lies on the line between metals and nonmetals, and this has led to some debate about its characteristics.

For example, the halogens get progressively darker as you move down the group - fluorine is a pale gas while iodine is a grey solid. Some chemists therefore predict that astatine is a dark grey-black. But others consider it more of a metal and predict that it is shiny, lustrous, and a semiconductor. In compounds, sometimes astatine behaves a little like iodine and sometimes a little like silver. For all of these reasons, it is often put to one side when discussing halogens.

Fig. 8 - The electron configuration of astatine

Tennessine

Tennessine is the final member of the halogens, but some don’t consider it a proper member at all. Tennessine has the atomic number 117 and is an artificial element, meaning that it is only created by colliding two smaller nuclei together. This forms a heavier nucleus that only lasts for a few milliseconds. Once again, this makes it just a little bit tricky to figure out!

Chemists predict that tennessine has a higher boiling point than the rest of the halogens, following the trend seen in the rest of the group, but that it doesn’t form negative anions. Most consider it to be a sort of post-transition metal instead of a true nonmetal. For this reason, we often exclude tennessine from group 7.

Fig. 9 - The electron configuration of tennessine

Reactions of group 7

The halogens take part in multiple different types of reaction, especially fluorine, which is one of the most reactive elements in the periodic table. Remember that reactivity falls as you go down the group.

Halogens can:

• Displace other halogens. A more reactive halogen will displace a less reactive halogen from an aqueous solution, meaning the more reactive halogen forms ions and the less reactive halogen is produced in its elemental form. For example, chlorine displaces iodide ions to form chloride ions and a grey solid, iodine.
• React with hydrogen. This forms a hydrogen halide.
• React with metals. This forms a metal halide salt.
• React with sodium hydroxide. This is an example of a disproportionation reaction. For example, reacting chlorine with sodium hydroxide produces sodium chloride, sodium chlorate, and water.
• React with alkanes, benzene and other organic molecules. For example, reacting chlorine gas with ethane in a free radical substitution reaction produces chloroethane.

Here is the equation for the displacement reaction between chlorine and iodide ions:

${\mathrm{Cl}}_2+2{\mathrm{I}}^{-}\to 2{\mathrm{Cl}}^{-}+{\mathrm{I}}_2$

Halide ions can also react with other substances. They can:

• React with sulfuric acid to form a range of products.
• React with silver nitrate solution to form insoluble silver salts. This is one way of testing for halides, as you’ll see below.
• In the case of hydrogen halides, dissolve in solution to form acids. Hydrogen chloride, bromide, and iodide form strong acids, whilst hydrogen fluoride forms a weak acid.

Testing for halides

To test for halides, we can carry out a simple test-tube reaction.

1. Dissolve a halide compound in solution.
2. Add a few drops of nitric acid. This reacts with any impurities that might give a false-positive result.
3. Add a few drops of silver nitrate solution and note down any observations.
4. To further test your compound, add ammonia solution. Once again, note down any observations.

With any luck you should get results a little like the following:

Fig. 10 - A table showing the results of testing for halides

The test works because adding silver nitrate to an aqueous solution of halide ions forms a silver halide. Silver chloride, bromide, and iodide are insoluble in water, and partially soluble if you add different concentrations of ammonia. This enables us to tell them apart.

Uses of halogens

The halogens have myriad different uses in everyday life. We’ve already looked at some above, but further examples include:

• Fluoride is an essential ion for animal health and helps strengthen teeth and bones. It is sometimes added to drinking water and you’ll commonly find it in toothpaste. The biggest industrial use of fluorine is in the nuclear power industry where it is used to fluorinate uranium tetrafluoride, ${\mathrm{UF}}_6$.
• Most chlorine is used to make further compounds. For example, 1,2-dichloroethane is used to make the plastic PVC. But chlorine also plays an important role in disinfection and sanitation.
• Bromine is used as a flame retardant and in some plastics.
• Iodine compounds are used as catalysts, dyes, and feed supplements.