Chemistry AS - Chapter 11 - The Halogens

Group 7 is on the right hand side of the periodic table and is made up on non-metals which exist as diatomic molecules (F2, Cl2, Br2, I2). They are called halogens.

The halogens vary in appearance. At room temperature:

• Fluorine is a pale yellow gas
• Chlorine is a greenish gas
• Bromine is a red-brown liquid
• Iodine is a black solid

They get darker and denser as we move down the group. They all have a "swimming-bath" smell.

A number of properties of fluorine are untypical. These stem from the fact that the F-F bond is unexpectedly weak compared with the trend for the rest of the halogens.

The small size of the fluorine atom leads to repulsion between non-bonding electrons because they are so close together.

The atoms get bigger as we go down a group because each element has one extra filled main level of electrons compared with the one above it.

As you move down the group, atomic radius, melting point and boiling point increase. As you move down the group, electonegativity decreases.

Electronegativity is a measure of the ability of an atom to attract electron density towards itself within a covalent bond.

Electronegativity depends on the attraction between the nucleus and bonding electrons in the outer shell. This in turn depends on the nuclear charge and the distance between the nucleus and the bonding electrons, plus the shielding effect of inner electrons.

The hydrogen halides. The shared electrons in the H-X bond get further away from the nucleus as the atoms get larger down the group. Shared electrons get further away from the halogen nucleus and increases shielding by more inner shells of electrons. These factors are more important than charge so electronegativity decreases as we go down a group.

Halogens usually react by gaining electrons to become negative ions with a charge of -1. These reactions are redox reactions and halogens are oxidising agents which are being reduced.

Cl2 +2e- -----> 2Cl-  (gain of electrons)

The oxidising ability of the halogens increase as we go up the group.

Fluorine is one of the most powerful oxidising agents known.

Halogens will react with metal halides in solution in such a way that the halide in the compound will be displaced by a more reactive halogen (not a less reactive one). 

This is called a displacement reaction.

Chlorine will displace bromide ions but iodine will not.

Cl2 (0) + 2NaBr (-1 on Br) ---> Br2 (0) + 2NaCl (-1 on Cl)

The two colourless materials react to produce the red-brown colour of bromine.

In this redox reaction the chlorine is acting as the oxidising agent by removing electrons from Br- and so oxidising 2Br- to Br2. The oxidation number increases from -1 to 0.

A halogen will always displace the ion of a halogen below it in the periodic table.

We cannot investigate the effect of fluorine in an aqeuous solution because it reacts with water.

Halide ions can act as reducing agents. They lose electrons to become halogen molecules. There is a definite trend in their reducing ability which is linked to the size of the ions.

The larger the ion, the more easily it loses an electron. This is because the electron is lost from the outer shell which is further from the nucleus as the ion gets bigger.

As you move down the group, the reducing power increases and their radius' increase.

This trend can be seen in the reactions of solid sodium halides with concentrated sulfuric acid.

Sodium solid halides all react with concentrated sulfuric acid. The products are different and reflect the reducing powers of the halide ions.

Sodium Chloride (Solid) - In this reaction, drops of sulfuric acid are added to solid sodium chloride. Steamy fumes of hydrogen chloride are seen. Product is sodium hydrogensulfate.

It is not a redox reaction as no oxidation state has changed. The chloride ion is too weak a reducing agent to reduce the sulfur (oxidation state +6) in sulfuric acid. It is an acid-base reaction.

Sodium Bromide (Solid) - In this case we see steamy fumes of hydrogen bromide and brown fumes of bromine. Colourless sulfur dioxide is also formed. Two reactions occur:

Sodium hydrogensulfate and hydrogen bromide are produced in an acid-base reaction.

NaBr +H2SO4 --> NaHSO4 + HBr

However, bromide ions are strong enough reducing agents to reduce sulfuric acid to sulfur dioxide. The oxidation state of the sulfur is reduced from +6 to +4 and bromine increases from -1 to 0. A redox reaction + exothermic.

Sodium Iodide (Solid) - Steamy fumes of hydrogen iodide and black solid of iodine is seen. Egg smell of hydrogen sulfide is also present. Yellow solid sulfur may also be seen. Colourless sulfur dioxide is also evolved.

Hydrogen iodide is produced in an acid-base reaction as before.

NaI(s) +H2SO4(l) ---> NaHSO4(s) +HI (g)

Iodide ions are better reducing agents than bromide ions so they reduce the sulfur in sulfuric acid even further (from +6 to 0 and -2) so that sulfur dioxide, sulfur and hydrogen sulfide gas are produced.

During the reduction from +6 to -2, the sulfur passes through oxidation state 0 and some yellow, solid sulfur may be seen.

All metal halides (except fluorides) react with silver ions in aqueous solution. Example: silver nitrate, to form a precipitate of the insoluble silver halide.

Cl-(aq) + Ag+(aq) ---> AgCl(s).

Silver fluoride does not form a precipitate because it is soluble in water.

Dilute nitric acid HNO3 or (H+ + NO3-) is first added to the halide solution to remove any soluble carbonate or hydroxide impurities as these would interfere with the test and form silver carbonate/hydroxide.

Add a few drops of silver nitrate solution and the halide precipitate forms. The reaction can be used to test for halides because we can tell from the colour of the precipitate which halide has been formed.

The colour of silver bromide and silver iodide are similar but if we add concentrated ammonia, silver bromide dissolves but silver iodide does not.

Silver fluoride - no precipitate

Silver chloride - white precipitate - dissolves in dilute ammonia

Silver bromide - cream precipitate - dissolves in concentrated ammonia

Silver iodide - pale yellow precipitate - insoluble in concentrated ammonia

Chlorine is a poisonous gas. However, it is soluble in water and in this form has become essential in treatment of water.

Chlorine reacts with water in a reversible reaction to form chloric (I) acid (HClO) and hydrochloric acid (HCl).

In this reaction, the oxidation number of one of the chlorine atoms (diatomic) increases from 0 to +1 and the other decreases from 0 to -1. This type of redox reaction, where the oxidation state of the same atoms of the same element increase and others decrease is called disproportionation.

This reaction takes place when chlorine is used to purify water for drinking and in swimming baths to prevent diseases.

Chloric acid is an oxidising agent and kills bacteria by oxidation. It is also a bleach.

Other halogens react similarly but much more slowly as you go down the group.

In sunlight a different reaction occurs which turns chlorine into hydrochloric acid and oxygen. Shallow pools need frequent addition of chlorine.

An alternative to direct chlorination of swimming pools is to add solid sodium (or calcium) chlorate (I). This dissolves into water to form chloric (I) acid in a reversible reaction.

In alkaline solution, the equilibrium moves to the left and the HClO is removed as ClO- ions. To prevent is happening, the swimming pools need to be kept slightly acidic. However, this is carefully monitored and the "water" never gets acidic enough to corrode metal components and affect swimmers.

Chlorine reacts with cold, dilute sodium hydroxide to form sodium chlorate which is an oxidising agent and active ingredient in bleach. This is also a disproportionation reaction.

Other halogens behave smiliarly.