kyro8008
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Strange stuff... Sulphuric Acid
Hi, I recently obtained some Sulphuric Acid (drain cleaner) and having never played with conc H2SO4 before preceded to play...
Out of curiosity I decided to see how vigorisly it reacted with metals. In two separate test tubes I added a strip of Al and Mg to the conc acid. It
foamed a little but then neither really seemed to do much, so I added a little water to each.
This considerably speeded things up. I was wondering why the acid didnt really do much when concentrated (it is Min.92% and I presume the rest is
mostly water), but did when water was added. Was it the burst of heat from the addition of water or having more water to ionise in?
Also I smelt bad eggs, and imediately tipped the contents down the sink followed by lots of water... But I dont know which or both tubes caused the
smell, I presume the Al might have been able to reduce the Sulphuric Acid, but this was far from expected.
I was expecting a displacment reaction and not H2SO4 being reduced to H2S.
If wanting to dissolve metals, is it better to dilute the acid first? and what about temp?
Thanks
[Edited on 27-8-2005 by kyro8008]
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woelen
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Concentrated sulphuric acid indeed does not react that vigorous with many metals. If it reacts with metals, then it usually oxidizes the metal, itself
being reduced to sulphur dioxide and water. For many metals, heating is required, before it reacts at an appreciable rate.
When the acid is diluted, then it reacts more easily with most metals, although not with all of them (copper and the more noble metals do not react
with dilute H2SO4). Many metals indeed reduce the sulfate ion SO4(2-) to S(2-) as a side reaction. So, most of the metal reacts according to the
standard acid/metal reaction, where H2 is formed, but a small fraction reduces the sulfate to sulfide You can also have formation of small amounts of
sulphur and of small amounts of sulphur dioxide. When you want to dissolve a metal for making a pure metal salt, then H2SO4 is not the best choice,
because of the side reactions, which give an impure product. The smell of H2S can be observed with almost every somewhat more electropositive metal,
such as zinc, magnesium, aluminium, the lanthanides.
If you want a pure metal salt, then dissolving the metal in aqueous HCl is much better. Excess acid can be driven off by heating and there is no side
reaction, in which the anion is reduced.
All oxo-acids suffer from these side reactions, e.g. H3PO4 can give rise to formation of PH3 and HNO3 usually gives nitrogen oxides, when a metal is
dissolved, but dilute HNO3 with zinc also results in formation of small amounts of ammonia.
Conc. H2SO4 is interesting for its dehydrating properties, its action on many organics in combination with conc. HNO3 and sometimes as oxidizer.
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Duster
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Bad eggs?
In firefighting we were taught that hydrogen sulfide smells of almonds, and once you stop smelling it the gas is in sufficent amounts to kill you (so
as long as you smell it your in the clear, ish)...
Unless I got that mixed up with something else, but im 90% sure H2S smells like almonds...
Damnit now this bugs me and Im gonna have to go check..
Duster
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garage chemist
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@ Duster: Either you or they got that mixed up.
H2S smells like rotten eggs, and yes they are right that it deadens the sense of smell as soon as the concentration is dangerously high.
The almond smell is HCN, but this doesn't deaden the sense of smell (but not everyone can smell it in the first place).
Concentrated H2SO4 reacts very different than dilute H2SO4. The oxidising property is very pronounced with the concentrated acid, and the acidic
properties are less pronounced due to lack of ion formation.
It can be reduced to even H2S if the metal is a strong enough reducer. But SO2 is mostly the main product of reduction.
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12AX7
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| Quote: | Originally posted by garage chemist
It can be reduced to even H2S if the metal is a strong enough reducer. But SO2 is mostly the main product of reduction. |
"Main product". Yep, quirky thing about H2S is you can smell a few ppb, so it could be a side reaction way the hell down on the
chemical's list of priorities and still make everyone run out of the room!
BTW I've noticed three colors of green vitriol: besides just green, I've seen the cyan photographed twice in its thread, and a colorless
form as an early product, when the liquor still stinks from those side products...any thoughts on that?
Tim
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ChemicalBlackArts
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The dilution of concentrated (or nearly so) acids is often a very exothermic process. The addition of the heat expedited the reaction rate
considerably.
--Chemistry in my veins--
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BromicAcid
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ChemicalBlackArts, while you were at it you could have restated the little rule of thumb "Always add acid to water never water to acid"
(although this mainly applies only to concentrated dehydrating acids such as phosphoric and sulfuric and perchloric [though sometimes dilluting that
can be dangerous]) if you do the reverse, expecially if only a little water is added, the water and acid can form two phases as the water is less
dense then the concentrated acid and boiling/popping can occur at the interface.
Also even if you do add the acid to the water it will heat up greatly just as ChemicalBlackArts stated and therefor if you are mixing in a plastic
container it should be dilluted in stages or better yet use some pyrex, depending on the quanity you are dilluting you may still want to do it in
stages.
Finaly, on your initial question on why the attack was not too vigorous on the aluminum and magnesium, in addition to the concentrated acid acting
slower for a number of reasons on its own, depending on the brand some of this may be due to corrosion inhibitors that are added, remember this stuff
is drain cleaner and they don't want it entirely decimating some of the pipes it comes into contact with.
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12AX7
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My drain cleaner is certainly exothermic with water, but I've never been able to get any splattering, and honestly I've tried  Not concentrated enough maybe? I don't recommend anyone taking my word for
it of course!
Tim
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