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Getting hydrogen from water

June 9, 2009

Chemistry is not my strong point; I generally found it to be arbitrary and hard to understand.  But it was full of dramatic demonstrations.  And if there’s one thing we love at Gravity and Levity, it’s science-themed drama.

And so, I thought I would share with the internet my favorite experiment from high school chemistry class.  It’s a pretty well-known one, but it’s worth retelling for anyone who hasn’t seen it before.  It’s also extremely easy to do at home (all you need is some salt and a 9-Volt battery), and if I can inspire a few people to perform a potentially dangerous experiment in their homes in the name of science, then I’ve done my job.

The theme of the experiment is getting hydrogen gas from water.  Before I first saw this experiment, I always thought of water as the world’s most stable compound.  It covers 70% of the earth and makes up 60% of the human body, and somehow in my mind that equated to it being perfectly “safe” and non-reactive.  But it turns out that splitting apart dihydrogen monoxide (water) is fairly simple.  All you need is a battery.

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A battery is a fairly complicated thing.  But there is a simple way to think about it that generally guides you in the right direction.  I always imagine it as a pair of reservoirs: one containing positive charge and another containing negative charge.  Something like this:

battery

These two “charge tanks” sit under the two terminals of the battery, and when the battery is connected to a circuit the negatives (electrons) flow from one tank to the other.  The voltage of the battery is a measure of how tightly-packed the charges are.  The closer the charges are to each other, the more strongly they will shoot out of their respective terminals when they are given a path to do so.  That’s why a battery’s voltage drops the more you use it: you are allowing the charge tanks to deplete, so that charges in a given tank don’t repel each other as strongly as they used to.  The lifetime of the battery is a measure of how big the tank is: it tells you how much current you can get out of the battery before it dies.  A 9V battery has a fairly high voltage (tightly-packed charges) but doesn’t last very long (the “tank” is small).  In contrast, one of those fat D battery has a low voltage (1.5 V) but can last a long time (it has a large tank).

So what happens if you put a battery in water?  Well, if you have pure water, not much.  The water molecule is electrically neutral, so it is not drawn to either terminal.  But if you mix some salt in the water, things are different.  Salt dissolves in water to leave behind positively-charged sodium ions and negatively-charged chlorine ions.  Once you put a battery in the water, the sodium ions migrate toward the “negative tank” and the chlorine ions migrate toward the “positive tank”.

At the positive tank, the chlorine ions get neutralized.  The chlorine ion Cl^- has an extra electron, and this electron gets ripped off and pulled into the positive tank.  The resulting neutral chlorine atoms then bond together (since they are more stable together) to form Cl_2, which is chlorine gas.  Chlorine gas is a serious poison, so be careful if you do this experiment at home.  Don’t do it on too large a scale (with your car battery or something), or it could be quite dangerous.

At the negative tank, something a little more complicated happens.  The electrons want to jump out of the tank, but they are more easily accepted by the water molecules than by the sodium ion (for some reason that is mysterious* to me).  As a result, the electrons jump to water molecules, which then become unstable and split apart.   The result is that each water molecule is split into one H (neutralized by the electron) and one OH^-.  Hydrogen atoms, like chlorine atoms, are more stable together, so two hydrogens get together and form H_2 gas.  The OH^- goes on to bond with the sodium left in the water or the iron on the battery leads (generally making the water dirty-looking).

The dramatic part, of course, is the hydrogen gas.  It bubbles quickly off the big lead of the battery, and you can collect it in a container if you do it carefully.  And since hydrogen is highly flammable, you can explode it.

I’m sure there are plenty of places on the internet where you can find better explanations of this experiment than I just gave.  Probably some of you reading this post could do better also.  But I highly encourage you to go out and have fun with this simple experiment.

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Here is a video made by Mrs. G&L and I, so you can see for yourself.  The explosion at the end is pretty weak, but you’ll get the idea.

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*Footnote

G&L reader Miriam points out that Na^+ doesn’t accept electrons for the same reason that it was so easily ionized in the first place: Na^+ has 10 electrons, which constitute two complete shells.  An additional electron would have to be placed in the third shell (because of the Pauli exclusion principle), which requires a significant amount of energy.

8 Comments leave one →
  1. Ed Stembridge permalink
    August 17, 2009 11:04 am

    I restored a 1950 Ford 8N tractor last year, and used an electrolysis tank to remove rust from all but the largest parts of the tractor. The tank was a large blow-up kiddie pool filled with water and washing soda (which is doing the same thing as the salt in your example). I submerged the metal part and placed a piece of rebar (metal) bent in a large circle surrounding the tractor part. I then connected the positive lead of a battery charger to the rebar and the negative lead to the part and turned it on. A day or so later, my part is rust and paint-free with only a coating of black oxide, which is easily removed with a pressure washer. I wipe a coat of phosphoric acid on the metal to stop it rusting again until I can get a coat of primer on it. I do this in a well-ventilated space because of the vigorous hydrogen production while it’s working. Google ‘electrolysis tank’ and you’ll get lots of hits for instructions on how to do this.

  2. Anonymous permalink
    August 27, 2009 6:05 pm

    Unfortunately, there is a mistake. If you drop a battery even in pure water, the electrolysis of water leaves you with oxygen at the anode and hydrogen at the cathode.
    And the chlorine in the salt added to the water does not significantly affect the process. About no chlorine gas will be formed because the reduction potential of water to oxygen is greater than that of the chloride ions.
    (Actually, putting hydrogen and oxygen in a 2:1 mole ratio gives you a better explosion🙂

    • Anonymous permalink
      April 16, 2011 12:32 pm

      Actually, there’s no mistake. The STANDARD reduction potential of water to oxygen is greater than that of the chloride ions, but once the cell starts to carry a current, conditions are no longer standard, and chloride ions are in fact reduced to chlorine gas. The brown color here is actually iron chloride, not iron oxide, or rust.

      • Anonymous permalink
        April 16, 2011 12:33 pm

        * The chloride ions are oxidized, not reduced

  3. Anonymous permalink
    February 14, 2012 9:59 pm

    So what happens to the sodium and oxygen?

  4. Jessaline permalink
    February 25, 2012 9:07 am

    thanks for he info i did this experiment and had sucess and told my science teacher and i really learned alot thanks🙂

  5. Edward permalink
    March 12, 2013 2:04 pm

    How does one measure the Hydrogen gas released? I’m doing an experiment similar to this and I need to know how much is released in 50mL of water,

  6. pa32r permalink
    June 20, 2015 12:52 am

    Oh, and kudos for Leo Kottke.

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