R06 Table of half-reaction equation equations for non-metals

At the top right-hand side of the Periodic Table we find the halogen group; fluorine, chlorine, bromine and iodine.
They are normally found as diatomic molecules, X₂.
The halogens (Greek for salt makers) are seldom found free in nature. This does not mean that the bonding between the atoms to form diatomic molecules is weak, but is rather a reflection of the fact that the halogens easily form ionic halides.

This is a redox reaction with a transfer of electrons taking place. Halogens are well known oxidising agents, which indicates that the halogen/halide ion redox couple has a potentiality for accepting electrons from another reagents i.e. oxidising that agent.

The potentiality of a particular X₂/2X^- redox couple relative to other X₂/2X^- redox couples can be qualitatively established by observing the effects of bringing different redox couples together. In illustration R6 two simple redox experiments are shown.

On the left-hand side of illustration R6 a brown liquid Br₂ is mixed with an aqueous solution of iodide ions for example from potassium iodide (KI). A light brown solution is obtained showing the presence of a small quantity of iodine. To demonstrate the presence of iodine more convincingly,a small volume of tetrachloromethane (CCl₄) is added and the vessel shaken. After leaving to stand for a few minutes the bottom layer of CCl₄ has a purple colour of I₂, while the aqueous layer has lost most of its brown colour. Another test would be to add a few drops of starch solution, whereupon the presence of a blueblack solution would indicate the presence of iodine. The total reaction which is occurring can be represented as:

In this reaction Br_2(aq) has oxidised to I_2(aq), while itself being reduced to . Br_2(aq) is therefore the oxidizing agent and the reducing agent.
The Br_2(aq)/2Br^- half- reaction equation is the stronger oxidiser and the two half-reaction equations can be written as :

In the reaction depicted on the right-hand side of illustration R6 chlorine (Cl₂), a light green gas,is bubbled through a colourless aqueous solution of sodium bromide (NaBr). A light brown aqueous solution of bromine results. Again the colour can be intensified by shaking with a small volume of tetrachloromethane. Since the bromine dissolves in CCl₄ better than in water, the CCl₄ layer at the bottom becomes brown. Bromine, like iodine, is apolar and therefore dissolves better in the apolar CCl₄ than in water which is highly polar.
The reaction can be represented by two half-reaction equation equations:

It is not possible by means of such a simple reaction to show that fluorine reacts as an oxidising agent with respect to chlorine, because it reacts violently with water in a redox reaction during which a gas, HF, is produced.
I₂, Br₂, and Cl₂, also react with water but less violently. We may assume that a redox reaction between F_2(aq) as oxidising agent and as the reducing agent can be represented by the half-reaction equations:

From these experiments it can be concluded that iodine is the least oxidising of the four halogen elements and that the oxidising power i.e. the potentiality of the X₂/2X^- couple to accept electrons increases in the order

iodine < bromine < chlorine < fluorine

Although these simple experiments are not all carried out under the standard conditions of temperature and pressure necessary for the measurements of standard reduction potentials, they achieve their purpose of demonstrating convincingly the order of their potentiality to accept electrons in the non-metallic halogen group.

N.B.: When working with any halogen element all necessary safety precautions should be taken.