Ask Ethan: What’s the quantum reason that sodium and water react?
Drop a chunk of sodium metal into water, and a violent reaction ensues. But it’s more than just chemistry at play.
“Chlorine is a deadly poison gas employed on European battlefields in World War I. Sodium is a corrosive metal which burns upon contact with water. Together they make a placid and unpoisonous material, table salt. Why each of these substances has the properties it does is a subject called chemistry.”
Sometimes, we learn things early on in life, and simply accept that’s how the world works. Drop a chunk of pure sodium in water, for example, and the reaction is legendary in its violence. As soon as you get that chunk of metal wet, the reaction fizzes and heats up, the sodium bounces around on the surface of the water, and even flames are produced. Sure, it’s just chemistry. But at a fundamental level, isn’t there something more going on? That’s what our reader Семен Стопкин (Semen Stopkin, from Russia) wants to know:
Which forces drive chemical reactions, and what takes place on a quantum level? In particular, what happens when water interacts with sodium? [Translated from Russian by physicist A. Vikman.]
The sodium/water reaction is a classic, and does have a deeper explanation. Let’s start by watching the reaction unfold.
The first thing you have to know about sodium is that, on an atomic level, it has just one more proton and one more electron than an inert, noble gas: neon. A noble gas is famed for not reacting with anything, and the reason is that all of its occupied atomic orbitals are completely full of electrons. That ultra-stable configuration gets ruined when you go up by one element on the periodic table, and this happens for all the elements that fit this pattern. Helium is ultra-stable, but lithium is highly reactive. Neon is stable, but sodium is reactive. And argon, krypton, and xenon are stable, but potassium, rubidium, and cesium are reactive.
The reason? It’s the extra electron.
When we learn about atoms, we learn to think about the nucleus as a hard, small, positively charged core at the center, and the electrons as negatively charged points that orbit it. But in quantum physics, that’s not really the whole story. Electrons can behave like points, particularly if you fire another high-energy particle or photon at them, but when left to their own devices, they spread out and behave like waves. Those waves can configure themselves in particular fashions: spherically (for the s-orbitals, which take 2 electrons each), perpendicularly (for the p-orbitals, which take 6 electrons each), and so on up through the d-orbitals (taking 10 electrons), the f-orbital (taking 14), and more.
The reasons these shells fill up is due to the Pauli Exclusion Principle, which prevents any two identical fermions (like electrons) from occupying the same quantum state. In an atom, if you have a full electron shell or orbital, the only place to put an additional one is in the next orbital up. An atom like chlorine will readily accept an additional electron, since it only requires one more to fill its electron shell; conversely, an atom like sodium will readily give up its last electron, since it has one extra electron over what will fill a shell. This is why sodium-chloride is such a good salt: the sodium gives up an electron to the chlorine, and both atoms are in a more energetically favorable configuration.
In fact, the amount of energy required for a neutral atom to give up its outermost electron, known as its first ionization energy, is especially low for all of those metals with one valence electron. If you look at the numbers, it’s much easier to strip a single electron off of lithium, sodium, potassium, rubidium, cesium, etc., than any other element.
So what happens in the presence of water? You might be tempted to think of water as its own very stable molecule: H2O, with two hydrogens bonded to one oxygen. But water is a highly polar molecule, meaning that one side of an H2O molecule (the side facing away from the two hydrogens) has a preferentially negative charge, while the opposite side has a preferentially positive charge. It’s a significant enough effect that it causes some water molecules — about one in a few million or so — to dissociate into two ions: a single proton (H+) and a hydroxyl ion (OH-).
This has a lot of consequences for things like acids and bases, dissolving salts, activating chemical reactions, etc. But the relevant one here happens when you add sodium. Sodium, this neutral atom with a loosely-held outermost electron, is now in the presence of water. This isn’t just the neutral H2O molecules, but the hydroxyl ions and the individual protons. The protons are the most relevant, and that leads to the key energy question we need to ask:
What’s more energetically favorable? Having a neutral sodium atom (Na) paired with a single proton (H+), or having a sodium ion that’s lost one electron (Na+) paired with a neutral hydrogen atom (H)?
The answer is a no-brainer; in pretty much every case, the electron will jump from the sodium atom to the first single proton it finds.
That’s why the reaction happens so quickly, and gives off so much energy. But the story isn’t complete. Now, you’ve made neutral hydrogen atoms, and unlike sodium, it doesn’t just form a block of individual atoms that you can bind together. Instead, hydrogen is a gas, and goes towards an even more energetically favorable state: forming the neutral hydrogen molecule, H2. So now, you’ve got lots of free energy (which goes into the heat of the surrounding molecules), neutral hydrogen gas, and it rises up out of the aqueous solution and into the atmosphere, which contains neutral oxygen gas (O2).
Get enough energy together, and the oxygen and hydrogen will react, too! This fiery combustion reaction produces water vapor, but also gives off even more energy. This explains why, when you drop a large enough chunk of sodium (or any group 1 element from the periodic table) into water, you get that tremendous, explosive release of energy. It’s all driven by the transfer of electrons, which occurs due to the quantum rules governing the Universe, and the electromagnetic properties of the charged particles that make up these atoms and ions.
So to recap, when you drop a chunk of sodium into water, here’s what happens:
- The sodium immediately gives up its outermost electron to the aqueous solution that is water,
- where its absorbed by a hydrogen ion, forming neutral hydrogen,
- with that initial reaction liberating a large amount of free energy, causing the surrounding molecules to heat up,
- then the neutral hydrogen becomes molecular hydrogen gas and rises out of the aqueous solution,
- and finally, if there’s enough energy, the atmosphere’s oxygen reacts with the hydrogen gas, creating a combustion reaction.
All of this can be explained simply and elegantly with the rules of chemistry, and that’s how it’s most often presented. Yet the very rules that govern the behavior of all of these chemical reactions arise from even more fundamental laws: those of quantum physics (like the Pauli exclusion rule, governing the behavior of electrons in atoms) and those of electromagnetism (governing how charged particles interact). Without those laws and forces, we’d have no chemistry at all! Yet thanks to them, any time you drop sodium into water, you know exactly what to expect. And if you haven’t learned your lesson yet, the answer is to wear protective equipment, don’t handle the sodium with your own hands, and stand back once the reaction occurs!
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Starts With A Bang is now on Forbes, and republished on Medium thanks to our Patreon supporters. Ethan has authored two books, Beyond The Galaxy, and Treknology: The Science of Star Trek from Tricorders to Warp Drive.