#melting points

Okay so, the general rule is the stronger the bonds the harder they are to break, therefore the more energy required to break them. This means the temperature will have to be higher before they can change into a gas as this involves them breaking away from each other a bit.

However the trends can’t just be explained from that alone because here are the first few elements

you need to know about all the other trends to explain it, because you need to know trends in bond strength.

Quick recap: 

note: remember the difference between a period and a group when talking about this, a period is a horizontal row while the group is the vertical column

In the second period, there is a gradual transition from relatively weak metallic bonding in lithium to strong network covalent bonding in carbon. Nitrogen, oxygen, and fluorine also form strong covalent bonds but they can’t form networks of bonds the way carbon does. Atoms of these elements pair up to form diatomic molecules. While the attractive forces within atoms in the diatomic gas molecules is strong, the forces between molecules is very weak. That causes the sharp dropoff in melting point after carbon in the second period.

Essentially, you need to know the different types of bonding in order to talk about the strength of these bonds

bonding in gases

here’s another graph yey!

All of the halogens exist as diatomic molecules - F2, Cl2, and so on. The intermolecular attractions between one molecule and its neighbours are van der Waals (dispersion) forces.

As the molecules get bigger there are more electrons which can move around and set up the temporary dipoles which create these attractions.

The stronger intermolecular attractions as the molecules get bigger means that you have to supply more heat energy to turn them into either a liquid or a gas - and so their melting and boiling points rise.

Metals

remember that metals work by forming lattices with delocalised seas of electrons. The more protons in the atoms, the stronger the positive charge and therefore the stronger the pull between the electron ‘sea’ and the nuclei, therefore the higher the melting/boiling point.

This only works when going along the period however because you’ve got another factor to take in - distance between electrons and nucleus. Once you start dramatically increasing the atomic radius, like going down a group, the melting point will fall because the electrons are further away from the positive centre. So: The melting and boiling points of the group 1 elements decrease on descending the group. This is due to a decrease in the forces of attraction between the atoms.

Giant molecular structures

These are the ones with covalent bonds, which are the strongest bonds so they have hella high melting points

this was actually a lot more difficult to explin than i’d thought, basically there are no trends to learn in non-gas groups. 7 & 8 both increase down the group. You will have to approach each question individually and look at the bonding in the elements

here’s another table for good luck

to summarize:

in gases: the bigger the molecule the stronger the van der Waals forces due to more electrons to set up the temporary dipoles

The more covalent bonds, the larger the melting point