The the outer electrons. All of this means

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The the outer electrons. All of this means

The halogens
in group 7 are always the oxidising agents to the halogens below. For example, fluorine
can oxidise chlorine, bromine, and iodine, taking an electron from either of
them to form its ion. They have then been oxidised and the fluorine has been
reduced. Chlorine can take electrons from bromine and iodine to form its ion, meaning
it can oxidise bromine and iodine and it is being reduced. Bromine will oxidise
iodine but not chlorine. And iodine doesn’t oxidise any of the above halogens.
This means that the oxidising abilities decrease as you go down the group. This
trend is largely to do with increased atomic radius as you go down the group,
as well as a decrease in nuclear attraction. Fluorine has a very small atomic
radius as it only has two shells. It also only has 9 electrons, so the strength
of the repel between them is very low compared to an element like iodine,
meaning the atom remains a smaller size without a lot of those forces pushing
the shells further out. It also doesn’t have many electron shells reducing the
strength of the nucleus on the outer electrons. All of this means that fluorine
has a very strong attraction to the atoms around it. When it is combined with
something with a weaker attraction such as chlorine, bromine or iodine, the
strong positive attraction from its nucleus can pull the outer electrons away
from the shell of these weaker atoms and therefore oxidise the other elements
and reducing itself. Chlorine is slightly larger than fluorine in atomic
radius. It also has more electrons and more shells. The increased amount of
electrons means there’s more negative charges repelling each other and
increasing the radius further. The increased amount of electromagnetic
shielding means there’s more interference between the nucleus and the outer
electrons, resulting in the nucleus having a weaker pull on its outer
electrons. This means that when combined with fluorine, chlorines positive
attraction isn’t strong enough to pull away the electrons on the fluorine atoms
as the positive charge of the fluorine’s atom is much stronger than the positive
charge coming from the chlorine atom. This is a similar trend going down the
group as the atomic radius increases and there’s more atomic shielding, the nuclear
attraction on the atoms outer electrons to get weaker. The elements can only
take electrons from the weaker ones below them as they don’t have the strength
to take them from the stronger positive charge of the smaller atoms above them.

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