# When you melt an ionic compound, do you "break" its electrostatic force of attraction, or its lattice energy?

I know that the melting point and the boiling point of ionic compounds are very high. However, when I was trying to find the reason for this, I found that this is because of the high electrostatic forces of attraction.

Does this mean that when we melt an ionic compound, we break its electrostatic forces? But then that means that in the liquid state, ionic compounds have only ions? That somehow doesn't seem right to me. I tried to find out if its the lattice of the ions that we are breaking instead but found no clear answer.

So, that is my question. When melting an ionic compound, do we break its lattice or its electrostatic forces? what about boiling?

Edit: I got this from Wikipedia:

Even when the local structure and bonding of an ionic solid is disrupted sufficiently to melt it, there are still strong long-range electrostatic forces of attraction holding the liquid together and preventing ions boiling to form a gas phase.[65] This means that even room temperature ionic liquids have low vapour pressures, and require substantially higher temperatures to boil.[65] Boiling points exhibit similar trends to melting points in terms of the size of ions and strength of other interactions.[65] When vapourized, the ions are still not freed of one another. For example, in the vapour phase sodium chloride exists as diatomic "molecules".[66]

Here, the last point. "Sodium chloride, in the gas phase, exists as diatomic molecules." This means, the electrostatic forces of attraction between the ions don't break ( The ions don't get separated). Like, for NaCl, you still have "molecules" of NaCl, you don't have ions of Na$$^+$$ and Cl$$^-$$ in the gas phase.

So, in a way, the only the lattice of NaCl broke, NaCl itself ( with the electrostatic forces) didn't break.

• And the difference between them is? Jul 21, 2020 at 17:33
• The term "lattice of NaCl" is misleading, possibly making false impression the lattice is formed by NaCl molecules. Jul 21, 2020 at 19:52
• No, your interpretation is still wrong. Covalent component of NaCl bonds is always present, as long as there's a bond. Jul 21, 2020 at 22:26
• What should "break its electrostatic forces" mean? As long as you have ions, there are electrostatic forces between them. There is nothing you can do about that, just as you cannot disable gravity between any two bodies.
– Karl
Jul 22, 2020 at 0:21

When you melt an ionic compound, do you "break" its electrostatic force of attraction, or its lattice energy?

When you melt an ionic compound, the ions are still there, and the electrostatic force of attraction is still there.

There is no longer a lattice. In the solid, there is a long-range order of the positions of the ions that you can describe by the translational symmetry of a repeating unit cell. This is gone in a liquid.

Lattice energy is not something on top of electrostatic potential energy. Instead, it is a way to describe these for a crystal structure.

Here, the last point. "Sodium chloride, in the gas phase, exists as diatomic molecules." This means, the electrostatic forces of attraction between the ions don't break ( The ions don't get separated). Like, for NaCl, you still have "molecules" of NaCl, you don't have ions of Na+ and Cl− in the gas phase. So, in a way, the only the lattice of NaCl broke, NaCl itself ( with the electrostatic forces) didn't break.

It is just the opposite: In a solid or liquid, there is no reason to pair up a cation with a specific anion. Each anion has several cations as closest neighbors, and each cation has several anions as closest neighbors (to be specific, 8 neighbors of the opposite charge in the LiCl crystal structure).

The difference between solids and liquid ionic substances is the number of neighbors while the distances don't change a lot. In lithium chloride, for example, the Li:Cl distance is 2.57 Å (and the Cl:Cl distance is 3.64 Å). In a simulation of molten lithium chloride, these distances are similar:

You can tell that the liquid LiCl is less densely packed by comparing the density of solid LiCl to that of liquid (or molten) LiCl, which is about 25% less (2.07 g/cm³ vs 1.502 g/cm³, source: http://moltensalt.org/references/static/downloads/pdf/element-salt-densities.pdf).

As Karl mentioned in his comment, the electrostatic forces are not being "broken" per se. The kinetic energy gained by the particles upon heating allows them to work against the electrostatic forces that are pulling them together, which results in a change of state (solid -> liquid -> gas). In simpler terms, the forces are present between the ions, but their effect is overcome. So, to answer your question:

Melting an ionic compound would break its lattice.