In this video, we will focus on Group VII elements, which we call the halogens. They are all very reactive non-metals. We call them the halogens because they react with most metals to form salts. Halogen is Greek for ‘salt-former’.
Group VII elements are called the halogens. The elements in this Group are fluorine, chlorine, bromine, iodine and astatine.
Halogens are non-metals that exist as molecules. Each molecule is made of two atoms covalently bonded together, hence we call them diatomic molecules.
The forces between the molecules are weak, little amount of energy is required to overcome this weak intermolecular forces of attraction. Hence, the melting point and boiling point of halogens are low.
Their boiling points increase down the Group. At room conditions, fluorine and chlorine exist as gases. Bromine exists as liquid, while iodine exists as solid.
All the halogens are coloured. The colour intensity increases down the Group. Fluorine is pale yellow, chlorine is pale yellow-green, bromine is reddish brown, while iodine is purplish black.
All halogens are poisonous and must be handled carefully in the Chemistry lab.
Formation of ionic compound
Halogens are very reactive non-metals, they react with metals to form ionic salts. For example, sodium burns in chlorine to form sodium chloride. A bright flame is observed in this reaction.
Halogens form an ion with a charge of -1. Each halogen atom has seven electrons in the valence shell. One electron is taken in to achieve stable noble gas octet electronic configuration. The ions of halogens are called halides. Ion of fluorine is fluoride, ion of chlorine is chloride, ion of bromine is bromide and that of iodine is iodide.
The compounds that halogens formed with metals are all ionic.
The halogens become less reactive down the Group. Fluorine is the most reactive non-metal in the Periodic Table. Chlorine is more reactive than bromine, and bromine is more reactive than iodine.
The interesting reaction halogens can undergo is displacement reaction. A more reactive halogen can displace a less reactive halogen from its halide solution.
For example, we have potassium iodide. When chlorine is added to potassium iodide, the more reactive chlorine will take the place of iodine, forming potassium chloride, leaving iodine back as an element.
Put it in chemistry terms, the more reactive chlorine displaces less reactive iodine from aqueous potassium iodide, forming potassium chloride and iodine.
Interesting observation can be made in this reaction – pale greenish yellow chlorine is bubbled into colourless solution of potassium iodide. A brown solution is formed due to the formation of iodine.
Cl2 + 2KI → 2KCl + I2
If we convert this chemical equation to ionic equation, we realised that this is a redox reaction.
Cl2 + 2K+ + 2I– → 2K+ + 2Cl– + I2
Cancel away the spectator ions, we will have Cl2 + 2I– → 2Cl– + I2. Here, you can see that the oxidation state of chlorine decreases from 0 in Cl2 to -1 in Cl–, while oxidation state of iodine increases from -1 in I– to 0 in I2. Hence, Cl2 has been reduced and I– has been oxidised. Chlorine is the oxidising agent while iodide is the reducing agent.
In a nutshell, halogens are non-metals and exist as diatomic covalent molecules. Their melting and boiling points are low and increases down the Group. These elements are coloured, and colour intensity increases down the Group.
They are all reactive, and reactivity decreases down the Group. The more reactive halogen displaces a less reactive halogen from its halide solution. All halogens are powerful oxidising agents.
Topic: The Periodic Table, O Level Chemistry, Singapore
If you would like to know more about the basics of Periodic Table, check out this post.
If you would like to know more about Group I elements, check out this post.
In the next post, we will focus on noble gases and transition elements.