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So as I've been told, when a substance melts, the actual bonds of the substance aren't broken, only the IMFs (inter-molecular forces). So why is it that metal groups decrease in melting points going down when they should be increasing because more shells exist thus ions in the lattice will be larger so more IMFs therefore it takes more energy to overcome. Instead, I'm being told that since the metallic bonds get weaker due to less attraction between electron sea and the positive ion, melting point decreases.

Could someone please help me figure out what I'm missing here?

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Depending whether you have a molecular, metallic or ionic compound, the independently moving particles in the liquid state are molecules, atoms and ions, respectively.

So for a molecular solid, you have to break the intermolecular forces to turn it into a liquid (or make them non-persistent so that the interaction partners can change over time).

For a metallic solid, you have to break the metal bonds intermittently so that the atoms can move in the liquid.

For an ionic solid, you have to break the ionic interactions intermittently so that the ions can move in the liquid.

Finally, for a network covalent solid, you would have to break covalent bonds to liquify it. For this reason, there is no liquid state of diamonds.

So why is it that metal groups decrease in melting points going down [...]

According to Jim Clark:

The strength of a metallic bond depends on three things:

  1. The number of electrons that become delocalized from the metal
  2. The charge of the cation (metal).
  3. The size of the cation.
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  • $\begingroup$ "no liquid state of diamonds"? chemistry.stackexchange.com/questions/6834/… $\endgroup$
    – Mithoron
    Feb 2 '20 at 17:37
  • $\begingroup$ @Mithoron Diamond describes a specific structure of a carbon allotrope; in contrast, NaCl describes a compound that has a distinct structure in the solid, liquid and gas state. So you can definitely obtain a liquid phase of carbon, but why call it diamond? $\endgroup$
    – Karsten Theis
    Feb 2 '20 at 18:35
  • $\begingroup$ Yeah, that's true. Even more, diamond turns into other phases of carbon before melting occurs (at least at reasonable pressure). $\endgroup$
    – Mithoron
    Feb 2 '20 at 18:40
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why is it that metal groups decrease in melting points going down

This isn't quite true for all metals. I think the only metals that behave like like this are the alkali metals. Group 6 metals increase in melting point. Calcium has a higher melting point than Magnesium. Etc.

more shells exist thus ions in the lattice will be larger so more IMFs therefore it takes more energy to overcome

I assume that shells here mean the electron shells. An atom having more electron shells usually causes it to have higher polarizability. And an atom that has higher polarizability has stronger dispersion forces. As a result, it will have higher melting point in comparison. Noble gases illustrate this. These compounds are inert and only interact with each other through weak intermolecular forces. Thus, radon has the highest melting point and neon has the lowest melting point. So you are right in this regard.

Metal atoms, however, share electrons with each other through something known as the metallic bond. Which is not an intermolecular force. A metallic bond is more like a covalent bond. There are a variety of factors that influence the strength of a metallic bond, as noted in Karten's answer.

the metallic bonds get weaker due to less attraction between electron sea and the positive ion

This isn't true for several metals either. Galium has a lower melting point than Indium. And many transition metals do not obey this trend either.

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  • $\begingroup$ It's imho abusing terminology to call interatomic or electron-nuclear attraction forces "molecular", but the attraction of an electron (shared or not) for neighboring nuclei is certainly what holds a material together. A metallic bond by any other name holds things together just as well. $\endgroup$
    – Buck Thorn ♦
    Feb 9 '20 at 10:58

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