Sodium Chloride vs Magnesium Oxide: Why MgO Melts So Much Higher

Sodium chloride melts at 801 °C. Magnesium oxide melts at roughly 2,852 °C. That is a difference of more than two thousand degrees between two white, crystalline, ionic solids that pack their ions in exactly the same way. If you were to build a model of each lattice from coloured balls, the two models would look identical apart from the sizes of the balls.

So why can one of them line the inside of a furnace, while the other would run off the walls as a liquid long before the furnace reached working temperature? This comparison works through the answer, which comes down to two numbers for each compound: the charge on the ions and the distance between them. It is one of the cleanest examples in chemistry of a property following directly from a simple physical law.

The two compounds at a glance

Feature	Sodium chloride, NaCl	Magnesium oxide, MgO
Ions	Na⁺ and Cl⁻	Mg²⁺ and O²⁻
Charge product (size only)	1 × 1 = 1	2 × 2 = 4
Formed from	Sodium (group 1) and chlorine (group 17)	Magnesium (group 2) and oxygen (group 16)
Lattice type	Rock-salt (face-centred cubic), 6:6 coordination	Rock-salt (face-centred cubic), 6:6 coordination
Approximate ionic radii	Na⁺ about 102 pm, Cl⁻ about 181 pm	Mg²⁺ about 72 pm, O²⁻ about 140 pm
Distance between neighbouring ion centres	about 282 pm	about 211 pm
Lattice enthalpy (magnitude)	about 787 kJ/mol	about 3,791 kJ/mol
Melting point	801 °C	about 2,852 °C
Solubility in water	Very soluble	Almost insoluble
Conducts electricity	Only when molten or dissolved	Only when molten
Typical use	Food, de-icing roads, feedstock for chlorine and sodium hydroxide	Refractory linings for furnaces and kilns, heat-resistant insulation

Lattice enthalpy values vary slightly between data books, because some are calculated from a Born–Haber cycle and some from a theoretical model, and different sources round differently. Treat the figures above as representative rather than exact. The ratio between them, close to 5, is what matters.

What stays the same: the rock-salt lattice

Both compounds crystallise in the rock-salt structure. Each positive ion sits at the centre of an octahedron of six negative ions, and each negative ion is surrounded by six positive ions in the same way. That is what “6:6 coordination” means. Beyond the six nearest neighbours of opposite charge come twelve next-nearest neighbours of the same charge, slightly further away, and so on outward through the whole crystal.

Because the geometry is identical, the arrangement contributes the same factor to the lattice energy in both cases. (Physicists call this the Madelung constant; for rock salt it is about 1.75.) This is useful for a comparison: it removes structure as a variable, like a fair test in which only the ions have been changed. If you want to see how other arrangements differ, such as caesium chloride or fluorite, the guide to ionic crystal structures walks through them.

What changes: charge

The attraction between two ions follows Coulomb’s law. The force, and the energy released when the ions come together, is proportional to the product of the two charges divided by the distance between their centres:

attraction ∝ (charge on cation × charge on anion) ÷ distance

In sodium chloride the charges are +1 and −1, so the product has a size of 1. In magnesium oxide they are +2 and −2, so the product is 4. With nothing else changed, every ion-to-ion attraction in MgO would be four times stronger than the matching one in NaCl.

This is the single biggest reason for the difference, and it is worth stating plainly for exams: doubling the charge on both ions multiplies the electrostatic attraction by about four, not by two. Students often write “the charges are doubled, so the bonding is twice as strong”, which undersells the effect.

What changes: size

The second factor is distance. An Mg²⁺ ion is much smaller than an Na⁺ ion. Both have the electron arrangement of neon, but magnesium has twelve protons pulling on those ten electrons instead of eleven, so its electron cloud is drawn in tighter. The oxide ion is also smaller than the chloride ion, because it has only two shells of electrons where chloride has three.

Add the radii together and the centre-to-centre distance falls from about 282 pm in NaCl to about 211 pm in MgO. Dividing by a smaller distance makes the attraction larger again, by a factor of about 282 ÷ 211, which is roughly 1.3.

Putting the two factors together

Multiply the two effects and the simple model predicts that MgO’s lattice energy should be about 4 × 1.3, or a little over 5, times that of NaCl. The measured ratio, 3,791 ÷ 787, is about 4.8. The small gap exists because real ions are not hard spheres: their electron clouds repel each other at close range, and that repulsion becomes relatively more important when ions are squeezed closer together. Even so, the back-of-an-envelope estimate lands close to reality, which is a good sign that the physical picture is right.

The detailed story of how lattice enthalpy is defined, measured and calculated is in the article on lattice energy.

From lattice enthalpy to melting point

Melting does not require the ions to be separated completely into a gas. It only requires the regular lattice to break down enough for ions to slide past one another. So lattice enthalpy and melting point are not the same quantity, and you cannot calculate one from the other with a simple formula. They are, however, strongly linked: a lattice in which every attraction is several times stronger needs far more thermal energy before the ions can shake loose from their positions.

That is why MgO stays solid at temperatures that would turn NaCl into a free-flowing liquid, and why magnesium oxide bricks are used to line furnaces. Sodium chloride would be useless there.

Why MgO forms at all

There is a puzzle hiding in the table. Making Mg²⁺ means removing two electrons from magnesium, and the second one costs a lot more energy than the first. Making O²⁻ is even odder: adding the first electron to an oxygen atom releases energy, but forcing a second electron onto an already negative ion costs energy, because the electron is repelled. Overall, forming a gaseous O²⁻ ion from an O atom is an endothermic step.

So why don’t magnesium and oxygen settle for Mg⁺ and O⁻ ions instead? Because the lattice enthalpy repays the extra cost many times over. A lattice of 2+ and 2− ions releases roughly four times more energy than a lattice of 1+ and 1− ions of similar size would. The expensive ionisation step is an investment that the lattice returns with a large profit. A Born–Haber cycle makes that energy bookkeeping visible step by step.

Other properties that follow

Solubility. Dissolving an ionic solid means pulling ions out of the lattice and surrounding them with water molecules. For NaCl the energy released by hydrating the ions is close enough to the lattice enthalpy that it dissolves readily. For MgO the lattice is so strongly held that water cannot pay the price, and it is almost insoluble. (It does react very slowly with water at the surface to give magnesium hydroxide, but that is a different process.)

Hardness. Both are brittle, like all ionic solids: shift one layer of ions sideways and like charges end up next to each other and push apart. MgO is harder, though, because its stronger attractions resist the initial deformation more.

Conductivity. Neither conducts as a solid, because the ions are locked in place. Both conduct when molten. NaCl also conducts in solution; MgO barely dissolves, so its solutions are not a practical test.

When each comparison matters

This pairing is a favourite exam question at the level where students first meet lattice enthalpy, because it isolates the charge effect so neatly. Use it when you need to:

• explain why ionic compounds with higher-charged ions have higher melting points;
• predict the order of lattice enthalpies in a set of compounds (for example NaF, NaCl, MgO, CaO);
• justify the choice of MgO as a refractory material;
• show why compounds of group 2 metals are often less soluble than those of group 1.

For a like-for-like comparison where only size changes, pair NaCl with NaBr or KCl instead. Charge is held constant, and the smaller effect of distance on its own becomes visible.

Common mistakes

• “MgO has stronger bonds because there are two bonds per ion.” There are no individual bonds between particular pairs of ions. Each ion is attracted to all its neighbours; the attraction is stronger because the charges are larger.
• “Doubling the charges doubles the attraction.” Both charges double, so the product goes up fourfold.
• “Mg²⁺ is smaller because it has fewer electrons than Na⁺.” Both ions have ten electrons. Mg²⁺ is smaller because it has one more proton pulling on the same number of electrons.
• “Melting breaks the ionic bonds completely.” Melting loosens the lattice so ions can move; full separation into gaseous ions is what lattice enthalpy describes.
• “MgO has a different structure, which explains the melting point.” The structures are the same. That is exactly what makes this a fair comparison.
• Forgetting the units or sign. Lattice enthalpy is often quoted as a positive value for breaking the lattice and a negative value for forming it. State which convention you are using.

Key takeaways

• NaCl and MgO share the same rock-salt lattice with 6:6 coordination, so the difference in their properties comes from the ions alone.
• The charge product rises from 1 to 4, which is the main reason MgO’s lattice enthalpy (about 3,791 kJ/mol) is nearly five times NaCl’s (about 787 kJ/mol).
• Smaller ions in MgO bring the charges closer and add a further boost of about a third.
• Stronger attractions explain MgO’s much higher melting point (about 2,852 °C against 801 °C), its greater hardness and its very low solubility.
• Forming 2+ and 2− ions costs extra energy, but the much larger lattice enthalpy more than pays it back.