Chlorine occurs in nature mainly as sodium chloride (table salt) in sea water and solid salt deposits, presumably from what was once salt lakes. The natural isotopes for chlorine are 35Cl (75.5%) and 37Cl (24.5%). Along with the chloride are bromide and iodide, in very much smaller quantities. Astatine, in Greek means unstable, is the product of the radioactive decay of uranium and thorium. We will ignore astatine in our discussion.
Chlorine is prepared industrially by the electrolysis of salt water (or brine).
This process will also work for hydrochloric acid, a by-product in chlorination process.
In the laboratory chlorine is prepared by the reaction;
HALIDES
The halides of the alkali metals and the alkaline earth metals are ionic, except for lithium and beryllium.
As the charge/radius ratio of the metal ions increases, covalence increases. Example, KCl (totally ionic), CaCl2, SeCl3, and TiCl4 - elements of the third row of the Periodic Table.
For a particular metallic element as the size of the halide ion increases, covalence increases. Example, AlF3 (ionic), AlCl3, AlBr3, AlI3 (covalent).
When the metals show more than one oxidation state, the lower halides tend to be ionic and the higher halide covalent. Example, PbCl2 (ionic) and PbCl4 (covalent).
All ionic halides are soluble in water, except silver (I), copper (I), mercury (I), and lead (II).
Tutorial 1
Explain why the solubility of the halides of the alkali metals and the alkaline earth metals are in the order: iodide > bromide > chloride > fluoride. Answer.
The hydrogen halides - HF, HCl, HBr, and HI - are considered covalent halides since they are gases at room temperature. However they dissociate completely into H+ and X‾ in water. The acid strength is in the order: HI > HBr > HCl. So the strength of the covalent bonds between two "highly unequal" partners is not strong.
Halogens can react with fluorine to give XF (except for IF), XF3, XF5, and XF7. Between the other halogens only BrCl, ICl, IBR, and ICl3 are known.
Iodine can also exist as a cation, I+, in solvents with lone electron pairs.
OXIDES
Each of the halogens can have several oxides. For example Cl2O, ClO2, Cl2O6 and Cl2O7. However they are very unstable, tending to explode. For example chlorine dioxide, ClO2, is a useful oxidising reagent, but because it is so unstable it is commercially prepared at the site and used immediately.
The most stable oxide is iodine pentoxide, I2O5. It is stable up to 300�C. It is the anhydride of iodic acid.
Iodine pentoxide is an oxidising reagent. It can react with carbon monoxide to give an iodine molecule.
This reaction is made used of to determine the amount of carbon monoxide, as the iodine produced can be easily determined.
The other oxo acid known is chlorous acid, HClO2. It is a weak acid (Ka = 10‾�) prepared by reacting barium chlorite with sulphuric acid, however it cannot be isolated in the free state.
The chlorites are generally prepared by dissolving chlorine dioxide, ClO2, in alkaline solutions.
These solutions are fairly stable and are used as bleaching agents.
Tutorial 2
- For the compound Cl2O, would you name it chlorine oxide or oxygen chloride? What if it is F2O?
- Why is barium chlorite used in the preparation of chlorous acid? Answer.
REACTION OF HALOGENS WITH BASES
Chlorine dissolves in a base to give a chloride ion and a hypochlorite ion.
When this solution is warmed the hypochlorite will disproportionate to a chloride ion and a chlorate ion.
The stability of the hypohalites are in the order; ClO‾ > BrO‾ > IO‾
IO‾ is so unstable that at room temperature it exist as a much of iodate and iodide.
PERACIDS
When solid potassium is heated it disproportionate to sodium chloride and sodium perchlorate.
Perchlorates of all metallic elements are known.
There are two interesting characteristics worth remembering. All perchlorates are soluble in water except potassium, rubidium and cesium. So far we know that all the alkali salts are soluble in water. The second is, the perchlorate anion, which is tetrahedral in structure, do not form complexes.
Both perchloric acid (HClO4) and perchlorates are not very reactive at room temperature, but at elevated temperature and high concentration it can react explosively. Perchloric acid is commercially available in 72% concentration.
There is no evidence of perbromate ion. This is in line with the general trend shown by the third-row anions - [GeO4]‾4, [AsO4]‾� and [SeO4]‾�. This is seen as the poor interaction between the 4d orbitals of the element with the 2p orbitals of oxygen.
Periodic acid (HClO4) and periodates resemble the perchloric acid and perchlorates, as good oxidising reagents, but with lower reactivity. So they found uses as analytical reagents for the quantitative determination of manganous ion.
Potassium periodate is a very stable compound at ambient condition and so is often used as a primary standard for the determination of thiosulphate conentrations.
Once the sodium thiosulphate is standardised it can be use to standardise iodine solutions. Iodine solutions are often used to determine quantitatively the amount of
antimony (III) in the presence of sodium bicarbonate. This is because the actual reaction is reversible and so the reaction is made to go to completion by removing the HI formed.
Sb2O3 + 2I2 + 2H2O � Sb2O5 + 4 HI HI + NaHCO3 � NaI + H2CO3 The carbonic acid is easily removed by warming the solution.
Tin (II) Sn+2 + I2 � Sn+4 + 2I‾
Sulphites [SO3]‾� + I2 + H2O � 2H+ + 2I‾ + [SO4]‾�
Other compounds that can be analysed quantitatively via iodine/thiosuplhate titration will be;
Hydrogen peroxide H2O2 + 2I‾ + 2H+2 � I2 + 2H2O
Iron (III) 2Fe+3 + 2I‾ � 2Fe+2 + I2
Tutorial 3
Given the following elements W, X, Y, and Z with atomic number of 16, 18, 19, and 20 respectively, which element forms an oxide that when dissolve in water gives a solution with pH less than 7? Give explanation. (From: Ms Vivien Li (15), Grade 10, Hong Kong) Answer.
