Natural states of elements down a group

[305920501]

I know a couple questions have shown up on exams about naturally occurring diatomic elements and their gaseous forms- I was wondering if there's any reason behind or trend to the states that elements prefer to be in. For example, down Group 17, I know F and Cl are diatomic and in a gas state at room temperature. Meanwhile, Br and I, both in the same group just below, are liquid and solid respectively at room temperature but also diatomic. Is there a reason for these changes down the group but but still a consistency of homonuclear diatomic molecules?

[Chem_Mod]

This trend is due to the increasing weight of the halogen as you go down the group! The iodine atom is massive compared to fluorine, and has many many more electrons. So when you have the molecule I₂, it is going to have pretty strong Van der Waals disperson forces keeping these molecules together. The molecule is very heavy and there are a lot of electrons that can "slosh" around and create temporary dipoles. So, I₂ is a solid. Br₂ is lighter than I₂, with weaker intermolecular forces, and will be a liquid. Cl₂ and F₂ are even lighter with even smaller intermolecular forces, so their molecules are free to spread out and make a gas! Since they have the same valence configuration, these atoms will bond in the same way, but their molecules exist in different states due to their differing weights.

I've moved this topic from coordination compounds to the intermolecular forces topic.