[Kemistry Kaiser]

Hey all,

*For context, I'm a sophomore chemistry student at Penn State so excuse my ignorance of advanced chemistry concepts.

Recently my friend and I have been extracting lithium from the core of AA energizer batteries and adding it to water. Long story short, it was a gateway reaction, and lead to us yearning for more of a show. I bought some HCl and misc. safety equipment from a hardware store and we've been using the acid instead of water which, as predicted, makes for a much more vigorous reaction. This led to us purchasing mineral oil to store the lithium under so we could drop a handful of semi-tarnished lithium metal into the acid all at once. When I did this, however, the solo cup of acid I added the lithium to simply began bellowing large quantities of  an unknown colorless gas, which continued for minutes, contrary to the quick, fiery reaction we were observing earlier. I'm assuming the mineral oil that was coating the lithium prevented it from igniting, but why the gas? The lithium didn't produce any noticeable gasses before under the same exact conditions without being stored under mineral oil, so what is the mineral oil doing to the reaction that's inducing the production of this gas? A professional opinion would be very much appreciated.

Also, is there any chance that the black iron disulfide (that acted as the cathodic strip in the batteries) will decompose into Dihydrogen sulfide or sulfur dioxide? I don't want to poison myself but at the same time would like to explore its properties. Thanks to anyone that responds!

[curiouscat]

You doing this indoor? 

[Arkcon]

Its hard to follow exactly what you're describing.  Let's try a different tack.  This is a learning forum, not a gee-wiz lookitwhaticandoooo blog.  You're a sophomore in college, so you can surely write out a high school level balanced equation.  Let's see that.  What reaction are you expecting with lithium and water, and lithium and weak acid?  What are you observing in those cases?  And what happens differently after a mineral oil dip?  Your mentioning a colorless gas really throws me -- what color gas did you get before?

[Kemistry Kaiser]

Sorry guys, I'm new to the site and I'm just trying to learn, and I do take serious qualitative notes during all of my backyard experiments. Here are my expected reactions and observations (again, a professional should verify these):

Lithium metal added to tap water:
Li(s)+2H₂O(l) Li⁺(aq)+2OH^-(aq)+H₂(g)

Observations: Immediate bubbling followed by small flames and eventually a sparkling explosion.

Lithium metal added to HCl:
Li(s)+2H3O⁺(aq) Li⁺(aq)+2H₂O(l)+H₂(g)

Observations: Almost instantaneously began producing flames and a sparkling explosion (significantly faster reaction than the Li metal added to tap water).

Lithium metal coated in mineral oil added to HCl: I'm really not entirely sure what happened after my observation of this test. In my notes I expected the reaction to be the same as Li added to HCl, but the has that bellowed out was so thick and was in no rush to rise, so again, I'm hesitant to claim it's hydrogen gas, which I would expect to rise quickly and not be as thick. I've never worked with gaseous hydrogen before so maybe I'm just ignorant. The only reason I make a point in saying the gas was colorless was because there was a lot of it and I didn't report any visible gas from the previous reactions. Now that I think about it, perhaps the hydrogen gas ignited and that contributed to the flames and sparks of the Lithium?

Observations: The only observations I reported for this reaction was very audible bubbling and bellowing colorless vapor.

NOTE: All reactions were carried out in polystyrene cups.

[Arkcon]

This is some good work, its rare to see someone in Citizen Chemist try to learn something, so +1 for you.  Now ...

Lithium metal added to tap water:
Li(s)+2H2O(l) Li+(aq)+2OH-(aq)+H2(g)

Observations: Immediate bubbling followed by small flames and eventually a sparkling explosion.

Lithium metal added to HCl:
Li(s)+2H3O+(aq) Li+(aq)+2H2O(l)+H2(g)

Why two different reactions?  You seem to realize that you're making the hydroxide of an active metal in the first case, why would you not get the same substance in the second case?

As for your different observations, you're seeing a different intensity of reactions, but maybe not a different product.  Except in the last case, but that reaction is very different, does it resemble anything else you've experienced outside of this reaction?

[AWK]

Lithium metal added to tap water:
Li(s)+2H₂O(l) Li⁺(aq)+2OH^-(aq)+H₂(g)


Lithium metal added to HCl:
Li(s)+2H3O⁺(aq) Li⁺(aq)+2H₂O(l)+H₂(g)

Reactions are wrongly balanced

[Kemistry Kaiser]

This is some good work, its rare to see someone in Citizen Chemist try to learn something, so +1 for you.  Now ...

Lithium metal added to tap water:
Li(s)+2H2O(l) Li+(aq)+2OH-(aq)+H2(g)

Observations: Immediate bubbling followed by small flames and eventually a sparkling explosion.

Lithium metal added to HCl:
Li(s)+2H3O+(aq) Li+(aq)+2H2O(l)+H2(g)

Why two different reactions?  You seem to realize that you're making the hydroxide of an active metal in the first case, why would you not get the same substance in the second case?

As for your different observations, you're seeing a different intensity of reactions, but maybe not a different product.  Except in the last case, but that reaction is very different, does it resemble anything else you've experienced outside of this reaction?

I wrote two because I figured the hydroxide produced, in the second equation, was in the presence of an acid and would therefor recombine to produce water. I'm sure the pH wouldn't be 7, but I figured a lot of water would be produced and the solution definitely wouldn't be as basic as the first equation. Please correct me if I'm wrong.

[Arkcon]

So you do realize that the second case is actually two reactions, the production of hydroxide and its neutralization.  And you should know they're both energetic.  So what questions still remain?

[Kemistry Kaiser]

None. I just wanted my thoughts to be checked over by a professional. Thanks Arkcon.

[Arkcon]

See, this is the thing I never understood.  We want to use metallic sodium, but its unsafe exposed to water or the moisture in the air, or it will produce hydrogen and possibly auto-ignite.  Fine.  I can follow that.  So what's the plan for safe storage?  Soak it in kerosene or, in your case, mineral oil.  Great, lets store something that auto-ignites in a bucket of fuel.

"Say, this package is ticking.  What should I do with it?"

"Put it in a cardboard box of oily rags until the bomb squad can get here"

I recognize that, before use, sodium is usually washed free of the kerosene or mineral oil with ether or other reaction compatible solvent.  Still, seems like an accident waiting to happen.

[pgk]

MAY BE IRRELEVANT BUT IT IS IMPORTANT:
Never store metallic sodium or lithium or potassium in dichloromethane (DMC) or in any other chlorinated or halogenated solvent. Also, never wash alkalies with halogenated solvents, due to unstable (metastable) carbenes that will be formed and thus, a SERIOUS EXPLOSION might occur at any moment. 

[Kemistry Kaiser]

See, this is the thing I never understood.  We want to use metallic sodium, but its unsafe exposed to water or the moisture in the air, or it will produce hydrogen and possibly auto-ignite.  Fine.  I can follow that.  So what's the plan for safe storage?  Soak it in kerosene or, in your case, mineral oil.  Great, lets store something that auto-ignites in a bucket of fuel.

"Say, this package is ticking.  What should I do with it?"

"Put it in a cardboard box of oily rags until the bomb squad can get here"

I recognize that, before use, sodium is usually washed free of the kerosene or mineral oil with ether or other reaction compatible solvent.  Still, seems like an accident waiting to happen.

Huh, I never thought about that. I guess storing all flammable liquids in one designated cabinet in the lab is a similar principle- Minimize chance of dangerous materials reacting by making an isolated storage-bomb out of them lol.

[Dan]

See, this is the thing I never understood.  We want to use metallic sodium, but its unsafe exposed to water or the moisture in the air, or it will produce hydrogen and possibly auto-ignite.  Fine.  I can follow that.  So what's the plan for safe storage?  Soak it in kerosene or, in your case, mineral oil.  Great, lets store something that auto-ignites in a bucket of fuel.

"Say, this package is ticking.  What should I do with it?"

"Put it in a cardboard box of oily rags until the bomb squad can get here"

I recognize that, before use, sodium is usually washed free of the kerosene or mineral oil with ether or other reaction compatible solvent.  Still, seems like an accident waiting to happen.

Huh, I never thought about that. I guess storing all flammable liquids in one designated cabinet in the lab is a similar principle- Minimize chance of dangerous materials reacting by making an isolated storage-bomb out of them lol.

This reminds me of something I found a few years ago at a previous workplace.

One day, we'd run out of a solvent I needed in the lab. It was something slightly unusual that wasn't stocked at our chemical stores, so I went to ask the group in the lab next door if they had any. They let me have a rummage around in their solvent cabinets and at the back of one of them I found an old glass bottle - probably from the 70s. The label was "anhydrous hexane (over Na)". Problem was, the coil of Na wire was protruding about 10 cm above the level of the solvent (presumably from decades of slow evaporation). Now that is dangerous. Amazingly the group in question didn't seem particularly worried about it... I didn't want to share a building with that bottle, so arranged for disposal myself.

[Kemistry Kaiser]

See, this is the thing I never understood.  We want to use metallic sodium, but its unsafe exposed to water or the moisture in the air, or it will produce hydrogen and possibly auto-ignite.  Fine.  I can follow that.  So what's the plan for safe storage?  Soak it in kerosene or, in your case, mineral oil.  Great, lets store something that auto-ignites in a bucket of fuel.

"Say, this package is ticking.  What should I do with it?"

"Put it in a cardboard box of oily rags until the bomb squad can get here"

I recognize that, before use, sodium is usually washed free of the kerosene or mineral oil with ether or other reaction compatible solvent.  Still, seems like an accident waiting to happen.

Huh, I never thought about that. I guess storing all flammable liquids in one designated cabinet in the lab is a similar principle- Minimize chance of dangerous materials reacting by making an isolated storage-bomb out of them lol.

This reminds me of something I found a few years ago at a previous workplace.

One day, we'd run out of a solvent I needed in the lab. It was something slightly unusual that wasn't stocked at our chemical stores, so I went to ask the group in the lab next door if they had any. They let me have a rummage around in their solvent cabinets and at the back of one of them I found an old glass bottle - probably from the 70s. The label was "anhydrous hexane (over Na)". Problem was, the coil of Na wire was protruding about 10 cm above the level of the solvent (presumably from decades of slow evaporation). Now that is dangerous. Amazingly the group in question didn't seem particularly worried about it... I didn't want to share a building with that bottle, so arranged for disposal myself.

That's a really interesting story. Makes you wonder if they even knew it was back there before you found it.