Group 17

• The halogens are a group of reactive non-metals that contrast strongly with the alkali and alkaline–earth elements which are reactive metals.
• The halogens are so reactive that they occur naturally only in compounds.
• The halogens can be obtained by oxidation of halide, Hal– ions.
• The oxidising ability of the halogens increases going up the group.
• Fluorine is one of the most powerful oxidising agents known.
• For example: Cl2 + 2e– gain of electrons 2Cl–.
• These reactions are redox reactions – halogens are oxidising agents and are themselves reduced.
• Halogens usually react by gaining electrons to become negative ions, with a charge of –1.
• All the halogens exist as diatomic molecules, X 2 , linked by a single covalent bond.
• The halogens are all volatile.
• Fluorine (F) has the electron configuration 1s 2 2s 2 2p 5.
• Chlorine (Cl) has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 5.
• Bromine (Br) has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 5.
• Iodine (I) has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4 d 10 5s 2 5p 5.
• Group 7, on the right-hand side of the Periodic Table, is made up of non-metals.
• As elements, the halogens exist as diatomic molecules, F 2 , Cl 2 , Br 2 , and I 2 , called the halogens.
• It is an acid–base reaction.
• Steamy fumes of hydrogen chloride are seen.
• The gaseous halogens vary in appearance, from pale yellow (fluorine) to dark brown (iodine), and have a characteristic ‘swimming-bath’ smell.
• There is a smell of rotten eggs because hydrogen sulfide is also produced: KI + H2SO4 → HI + KHSO4.
• The solid product is sodium hydrogensulfate.
• Hydrogen chloride is not further oxidised by concentrated sulfuric acid.
• Sodium chloride (solid) in this reaction, drops of concentrated sulfuric acid are added to solid sodium chloride.
• Sodium bromide (solid) in this case you will see steamy fumes of hydrogen bromide and brown fumes of bromine.
• The equations for the reaction with potassium bromide are: KBr + H2SO4 → HBr + KHSO4 2HBr + H2SO4 → Br2 + SO2 + 2H2O.
• With iodides, some steamy fumes of hydrogen iodide are formed, but the more obvious observation is a cloud of violet iodine vapour, mixed with some yellow sulfur.
• Colourless sulfur dioxide is also formed.
• Solid sodium halides react with concentrated sulfuric acid.
• The products are different and reflect the reducing powers of the halide ions shown above.
• This reaction is not a redox reaction because no oxidation state has changed.
• Sulfur dioxide is also produced but it is not visible because it is a colourless gas.
• In these reactions the halide ions lose (give away) electrons and become halogen molecules.
• With bromides, steamy fumes of hydrogen bromide are produced together with some orange gaseous bromine.
• Halide ions can act as reducing agents.
• This trend can be seen in the reactions of solid sodium halides with concentrated sulfuric acid.
• The larger the ion, the more easily it loses an electron.
• The chloride ion is too weak a reducing agent to reduce the sulfur (oxidation state = +6) in sulfuric acid.
• There is a definite trend in the reducing ability of halide ions, which is linked to the size of the ions.
• Many of the properties of fluorine are untypical, stemming from the fact that the F—F bond is unexpectedly weak, compared with the trend for the rest of the halogens.
• Halogens usually react by gaining electrons to become negative ions, with a charge of –1.