Is the melting and boiling point of ionic bond usually higher than covalent bond?

I know that compounds with ionic bonds are usually solid at room temperature, so I want other answers than this. (and this question can also be about why compounds with ionic bonds are usually solid in room temperature.)

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    $\begingroup$ Neither ionic nor covalent bonds have a "melting point". The melting point is a macroscopic property of a compund or element, while bonds are phenomena on the mircoscopic (atomic scale) level. Did you mean melting points of compounds incorporating ionic bonds as opposed to compounds incorporating solely covalent bonds? $\endgroup$ Commented Jun 21, 2013 at 14:31
  • $\begingroup$ @TanithRosenbaum - your comment is a good starting point for an answer that distinguishes between interatomic (intramolecular) forces and intermolecular forces, which govern the macroscopic properties. $\endgroup$
    – Ben Norris
    Commented Jul 22, 2013 at 12:49
  • $\begingroup$ In case of ionic compounds, there is a complete transfer of electrons from one atom of the element to another atom of the another element. So, there exists a strong electrostatic force. In case of covalent compounds, electrons are not completely dragged, the atoms of different elements are attracted by a weaker electrostatic force. Thus, ionic compounds have atoms bound by stronger electrostatic force than in case of covalent compounds. So, in general it is easy to break covalent bond than ionic. Thus, melting and boiling points of ionic compounds is greater than covalent compounds. $\endgroup$
    – Sensebe
    Commented Dec 6, 2013 at 13:06
  • $\begingroup$ The melting point of diamond (under pressure) is not lower than typical ionic compound. I guess one need to phrase the word "usually" in certain sense.... $\endgroup$
    – user26143
    Commented Dec 6, 2013 at 14:54
  • 3
    $\begingroup$ He should have said discrete covalently bonded whereby the actual bonding that defines the phase is physical or Van Der Waals. All network covalent molecules (especially 3D) have high melting points. $\endgroup$ Commented Dec 14, 2013 at 9:18

3 Answers 3


The answer relates to the strength of the interactions between the component units that make up a crystal or a solid.

The reason why anything is a solid at a given temperature is, crudely, that the interactions between the units that make up the solid (atoms, ions or molecules) are stronger than the amount of thermal energy available at that temperature.

Diamond, for example, is a solid because each unit carbon is bonded to its neighbours by a strong carbon-carbon bond that takes a great deal of energy to break. Non-polar molecular compounds, like candle wax (in reality a mixture but imagine a pure long chain hydrocarbon to keep it simple) are held together by the intermolecular forces between the molecules (often called dispersion or Van der Waals forces). These are much weaker than chemical bonds and depend on the surface area of the molecules. So small non-polar or not-very polar molecules tend to have low melting points. Hexane is a volatile liquid. The shape and symmetry of the molecule matters a bit to these interactions so toluene is a liquid down to -95°C but benzene, which is very similar to toluene but far more symmetric and slightly flatter only melts at about 6°C as the molecules pack together better encouraging stronger interactions. Polar molecules have stronger interactions because the molecular dipoles interact and this produces stronger forces than the non-polar dispersion forces. So phenol (very similar in size to toluene but far more polar) melts at about 41°C.

Simple ionic solid are made from charged ions. Table salt crystals consist of Na+ and Cl- ions. The key interaction between these is electrostatic and this is a strong force that takes a lot of energy to break. The melting point is about 800°C. Permanent electrostatic interactions are strong especially if the components can get close (remember, it's an inverse square law). This explains why not every ionic compound is a solid. There are ionic compounds that are liquid at room temperature. These tend to consist of molecular ions that are large and sometime irregular so forcing structures where the ionic force has to act over much larger distances and, therefore, be much weaker. An example is [BMIM]PF6 which has a melting point well below room temperature:


These ionic liquids are an interesting area of research in chemistry as they sometimes enable much less wasteful chemical reactions in industry (see wikipedia).


One set of comparisons is:

Why it take so much energy to turn sodium chloride molten, whereas candle wax melts virtually upon touch?

"Ionic compounds" tend to form crystals. You may have seen pictures of a crystal lattice of sodium chloride: cubic packing with a sodium being surrounded by many chlorides. It takes a lot of energy to break up all of those bonds.

In the case of wax (also a solid at room temperature), there are not these strong, "ionic" interactions between molecules. Things are more "layered" and disordered rather than crystalline. (Though there is definitely inter-molecular bonding Otherwise it would never be solid!!!)


It looks to me like sodium chloride is sort of a big frat party. All the boys are attracted to all the girls (you choose which element is which), so they're all holding hands with each other. To break NaCl into smaller bits, you have to break those handholds (melting, requiring lots of energy) or give them something else to hold (such as one end of a polarized water molecule). By contrast, sugar molecules are a bunch of close knit family groups. The members of each family are holding on to each other, but not attracted to any other family. To melt sugar, by getting the families to move around relative to each other, does not require much energy. When sugar is dissolved in water, the water molecules easily fit in among the sugar molecules, but again, elements of a given family are not attracted to outsiders.

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    $\begingroup$ Although the analogy is nice, it neglects much of the actual chemistry behind this effect. Perhaps you could go into more detail about why ionic and covalent compounds have these properties. $\endgroup$
    – bon
    Commented Nov 29, 2015 at 10:49

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