# Why is disodium oxide so unstable?

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1. Oct 30, 2016

### Tiiba

A powerful oxidizing agent, a powerful reducing agent. And yet, they form a relatively weak bond. Why?

2. Oct 30, 2016

### TeethWhitener

Wikipedia lists the standard enthalpy of formation as -416 kJ/mol. This is comparable to sodium chloride. What makes you think sodium oxide is unstable?

3. Oct 30, 2016

### Tiiba

The fact that this enthalpy of formation is from three atoms, while sodium chloride has two. It reacts with water. And sodium hydroxide dissociates into sodium and hydroxide, even though hydrogen is more electronegative.

4. Oct 30, 2016

### TeethWhitener

Even if you go by enthalpy of formation per atom, Na2O is more stable than water, with ΔHf = -286 kJ/mol. It's unclear what you mean by "unstable." Do you simply mean that Na2O undergoes reactions? In that case, you'd have to say that water is unstable as well, since a mole of water reacts with a mole of sodium oxide to give 2 moles of sodium hydroxide. The fact is that most compounds will react if you put them in the right environment. In fact, I'd wager that Na2O is probably the most stable stoichiometry of sodium and oxygen that you can make.

5. Oct 30, 2016

### Tiiba

Well, obviously, it's not a contact explosive. But when you put NaOH in water, you get Na+ and OH-. Hydrogen is more electronegative than sodium, yet oxygen prefers the company of hydrogen.

6. Oct 30, 2016

### Bystander

Dissociation is a property of strong electrolytes in water.

7. Oct 31, 2016

### Tiiba

8. Oct 31, 2016

### Staff: Mentor

I am not convinced by the data from the table you are right. It lists bond dissociation energy for the NaO dissociation - not for Na2O. While I can imagine NaO being present in gaseous phase at very low pressures, I am not sure its dissociation represents what is going on in Na2O.

9. Oct 31, 2016

### TeethWhitener

What's confusing me is that you keep bringing NaOH into a discussion about the stability of Na2O. It seems to me that you think that because Na2O reacts with water, and possibly also because the products of the reaction are soluble in water, this means that Na2O is unusually unstable, given oxygen's placement on the periodic table (maybe?). Is that the idea you're going for?

10. Oct 31, 2016

### Tiiba

I wanted to know why, in general, alkali metals prefer to bond with less electronegative elements like iodine over oxygen. Everybody here seems to think it's an illusion. Maybe it is.

11. Oct 31, 2016

### Bystander

Sodium iodide reacts with free oxygen, in solution or as a solid. What are you talking about?

12. Oct 31, 2016

### TeethWhitener

Ok. I think if you're basing your conclusions on the links you provided, Borek gave you the best answer: namely, that bond dissociation energies are measured for gas phase species and bear little relation to lattice energies that represent the forces holding ionic solids together. If you're looking at the comparative stability of sodium oxide and (say) sodium iodide, then Bystander gave you the best answer above. Based on the reaction:
$$2Na_2 O + 2I_2 \rightleftharpoons 4NaI + O_2$$
the equilibrium is to the left, meaning that Na2O beats NaI in terms of energy.
In terms of the reaction of Na2O with water, both species react to give NaOH. This has to do with the fact that 1) Na+ is very strongly solvated in water, and 2) O2- is extremely basic, to the point where it will immediately abstract a proton from water, giving OH-.

13. Oct 31, 2016

### Kevin McHugh

What are you talking about? The halides are the most electronegative elements in the periodic chart.

14. Oct 31, 2016

### Tiiba

As you can see from the chart I linked, iodine, and even chlorine, is beaten by oxygen. Halides are more electronegative than chalcogens of the same period, though.

So, I guess that in the gas phase, the bonds are weaker because the atoms feel the attraction only from their molecule, while in a crystal, they feel the attraction (and repulsion) of every atom that surrounds them. And somehow, even with the repulsion, that results in lattice energy being much higher than bond energy.

Never thought of that. I thought they're reporting the energy in the crystal. It's supposed to be at STP.

And I also guess that, despite lower electronegativity, chloride is more stable than oxide because it can form two weak bonds in place of one strong bond.

15. Oct 31, 2016

### TeethWhitener

Again, you're basing this conclusion on the gas phase stoichiometry. If you look at the crystal structure of solid phase Na2O vs NaCl, you see that a unit cell of NaCl has 4 Na and 4 Cl ions (rock salt structure), whereas sodium oxide has a unit cell with 4 Na and 2 oxygen ions (antifluorite structure). So if you go by the actual unit cell of the species, sodium oxide is more stable ion for ion than sodium chloride. In reality, a better metric for relative stability is the one Bystander pointed out above (which ends up falling pretty much in line with the enthalpies of formation that I mentioned initially).

16. Nov 1, 2016

### Tiiba

Wait, does sodium chloride burn in air? Because I don't think it does. For that matter, Wiki says NaI is non-flammable.

17. Nov 1, 2016

### Staff: Mentor

It doesn't.

18. Nov 1, 2016

### TeethWhitener

You also have to ask if sodium oxide reacts with chlorine gas (it doesn't). Which implies that the two compounds are roughly equally stable, as the enthalpies of formation tell you.

Also, sodium iodide yellows on exposure to oxygen and light, an indication that the I- is being oxidized to I2. Reaction with oxygen doesn't always have to cause a fire.

19. Nov 8, 2016

### DrDu

Nope, a primitive unit cell of NaCl contains only two atoms, 1 Cl and 1 Na. The cell you have in mind is not primitive.

20. Nov 8, 2016

### DrDu

I am here with Borek. Have a look at the following site:
https://usflearn.instructure.com/courses/986898/files/31116966
There is a list of lattice energies included and it shows that the sodiumoxide lattice has a much higher energy than that of the halides.