# Why does the melting point get lower going down the Alkali Metal Group with increase in atomic number?

Why does the melting point get lower going down the Alkali Metal Group with increase in atomic number?

You have to understand that melting point is related to the bonds in a metal between the metal cation and the 'sea of electrons'. With an increase in atomic number, you have an increase in electron shells.

As the radius of atoms get larger down the group, you could say that the force holding them together is spread over the greater area and hence, the metal cations are more weakly bonded. And coming back to my first point, when the bonding is weaker, the metal's melting point will decrease.

As you go down any group, there are extra electron shells (for example Lithium has 2 shells, Francium has 7) which cause electron shielding. Because the force of attraction must extend further and through these shells to the outermost shell, the force becomes less so there is less strong bonding occuring.

Thus, it is easier to break these less strong bonds, and because the melting point is a measure of the point at which all the bonds holding a metallic structure together are broken, it will decrease because this point is reached a lot sooner as the bonds are not as strong.

melting points of all group (i) elements is dependent on the strength of the metallic bond. In metallic bonding, the group (i) cations in the metallic lattice are attracted to the delocalised electrons. Down the group, the number of delocalised electrons and the charge on each cation remains the same at +1 but the cationic radius increases so the attraction between the cations and the electrons in the lattice get weaker down the group and so does the strength of the metallic bond

Why does the melting point get lower going down the Alkali Metal Group with increase in atomic number?

Ok, you're looking for a generalization of what happens to the melting point as you go down a column in the periodic table.

Think of the metal atoms as cations $\ce{M^+}$ and anions $\ce{M^-}$ packed into a crystal like NaCl. However for the metals think of the metal atoms switching + and - charges at a high frequency.

So as the atoms get bigger, the energy that is holding the crystal together gets smaller because the "ions" are further apart.

Note that this simple explanation doesn't explain "real" chemical bonding. So there are other layers of truth under this simple model.

1. Shell structure of an atom
2. Different crystalline lattices and hence energy "within" the crystal
3. Ionic , covalent, and metallic bonding

So the "real" truth is gleamed from the interactions of simple models into a more complex system.

You have to know that elements with high lattice energy has high melting and boilng point.And since lattice energy is inversely related with radius of atoms it increases down a group and hence the melting point increases. I hope this will help you!

• This doesn't follow. Atomic radius increases down the group so according to you lattice energy should decrease and so melting point should decrease. – bon Feb 10 '16 at 13:06
• The alkali metals are not ionic solids. – orthocresol Feb 10 '16 at 13:12

## protected by Martin - マーチン♦Feb 10 '16 at 16:50

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