## Covalent versus ionic bonds

Soon after I stopped blogging last year, I realized that I had a lot of things I wanted to say, just as before, and I had deprived myself of a channel for them. So I did the next best thing: I published notes on Facebook. For the first few days of this new blog’s life, I will simply cross-post those notes here, so if you’re already my friend on Facebook, you won’t be seeing any new material for a while; but I’m going to get through these one per day, so you won’t have long to wait.

We all know that covalent and ionic bonds are stronger than hydrogen bonds, dipole-dipole forces, and van der Waals forces (also known as London dispersion forces). But some textbooks go further and claim that covalent bonds are stronger than ionic bonds, or vice versa. Is this justified?

First of all, I would like to point out that ionic bonds generally become very weak in aqueous solution. This is because water is a highly polar solvent, that is, a medium with a high dielectric constant (about 82). This implies that the attractive force between two oppositely charged ions in water is 82 times weaker than the attractive force between the same two ions in vacuum, with all other things being equal. Thus, in aqueous solution, bonds with heavily ionic character, as in NaCl, become extremely weak, and nearly complete dissociation results; polar covalent bonds, as in carboxylic acids, are also weakened, though not to the same extent, resulting in partial dissocation; and nonpolar covalent bonds, as in hydrocarbons, are largely unaffected.

In the absence of a highly polar solvent, what would we be comparing if we compared the strengths of covalent and ionic bonds?

I can think of two approaches to this. In the first, we will consider melting and boiling points, which are bulk properties.

Ionic compounds are held in the condensed state by ionic bonding, whereas molecular compounds are held in the condensed state by dipole-dipole forces, hydrogen bonding, and van der Waals forces. Ionic compounds tend to have much higher melting and boiling points than molecular compounds. Therefore, we are justified in saying that ionic bonding is stronger than dipole-dipole forces, hydrogen bonding, and van der Waals forces.

There are also solids held together largely by covalent bonding. Examples include silicon and boron. These also have much higher melting and boiling points than molecular compounds. Therefore, we are justified in saying that covalent bonding is stronger than dipole-dipole forces, hydrogen bonding, and van der Waals forces.

However, when we try to compare ionic solids and network covalent solids, we can’t draw any definite trend in melting points. For example, the melting point of sodium chloride is 1074 K, whereas the melting point of calcium oxide is 2845 K. Meanwhile, the melting point of silicon is 1687 K, and that of boron is 2349 K. Data on boiling points is less complete in the literature; but we have silicon dioxide boiling around 2500 K and boron around 4200 K; sodium chloride boils around 1700 K, aluminium oxide around 3250 K, and magnesium oxide around 3900 K; and again no real trend can be seen.

In the second approach, we will attempt to compare individual bonds. Nonpolar covalent bond energies range from about 145 kJ/mol (oxygen-oxygen single bond) to about 945 kJ/mol (in dinitrogen). Typically encountered polar covalent bond energies range from about 305 kJ/mol (carbon-nitrogen single bond) to about 1070 kJ/mol (in carbon monoxide). Meanwhile, you will not find ionic bond energies tabulated anywhere, because there are no “ionic molecules”. When liquid sodium chloride boils, we do not obtain NaCl molecules, but rather Na atoms and Cl atoms; when liquid magnesium oxide boils, we find magnesium atoms and oxygen atoms, rather than MgO molecules.

So all we can do is attempt to estimate some figures. The Na-Cl distance in sodium chloride is about 282 pm, and we can calculate the electrostatic potential energy as $-k_C e^2/r \approx -8.2\cdot 10^{-19}\text{ J}$, or -490 kJ/mol. The Al-O distance in aluminium oxide should be about 190 ppm (estimated from ionic radii, since I can’t find the value online), which gives us $-k_C (2e)(3e) / r$ = 4400 kJ/mol; but these values don’t make too much sense, since the aluminium ion doesn’t carry a full 3+ charge, nor the oxide ion a full 2- charge. We can also try to obtain values using the lattice energies; for sodium chloride, the lattice energy is about (-) 790 kJ/mol, and each sodium is surrounded by six chloride ions, so we have six moles of bonds per mole of sodium chloride, giving the bond energy as about 130 kJ/mol; in aluminium oxide, the lattice energy is about -16000 kJ/mol (?), and each aluminium ion is surrounded by six oxide ions, so a mole of aluminium oxide contains twelve moles of bonds (since there are two aluminium ions per formula unit); this gives a bond energy of about 1300 kJ/mol. But it’s still questionable whether these values are meaningful, because we haven’t discussed the effect of repulsive forces that are also present in ionic crystals, and it’s not clear what should be done about them.

But to summarize, we find a lack of justification for calling covalent bonds or ionic bonds the stronger.