Craig, your comment on firefighting reminded me of a story. We once were in a meeting with some of our local firefighters and during a break I was chatting with one and asked how they evaluated the chemical hazards in a lab if they have
to enter. To paraphrase, his reply was "unless they're a hazmat officer, the average firefighter knows there are two kinds of chemicals: those you can spray with water and those you can't." I'm sure he was exaggerating, but I'm guessing broadly speaking it's
a fair assessment.
My impression is that a lot of how water reactivity is defined revolves around that: if you spray water at it during a fire, does this improve the situation (or at least have no effect), or does it make the situation worse? To me sulfuric
acid falls into the former category.
Sincerely,
Jack Reidy (he/him)
Research Safety Specialist, Assistant Chemical Hygiene Officer
Environmental Health & Safety
Stanford University
484 Oak Road, Stanford, CA, 94305
Tel: (650) 497-7614
From: ACS Division of Chemical Health and Safety <DCHAS-L**At_Symbol_Here**Princeton.EDU>
On Behalf Of Craig Merlic
Sent: Friday, February 18, 2022 12:28 PM
To: DCHAS-L**At_Symbol_Here**Princeton.EDU
Subject: Re: [DCHAS-L] water reactivity of sulfuric acid solution
The Chemistry here needs some correcting.
Bisulfate (HSO4-) is still a good acid as solutions have a pH of about 1. So it is corrosive and sodium bisulfate is used as "pool acid".
An "equimolar" solution of water and sulfuric acid is about 85% so is still intensely acidic and extremely corrosive.
Concentrated sulfuric acid is 98% and most chemists would say "dilute" sulfuric acid is 10%, although it is still acidic and not something to leave on your skin.
For storage, Richards comments are appropriate.
For "firefighting" let's bring in some logic. First it does not catch fire! Second, concentrated sulfuric acid does exothermically mix with water so splattering is an issue.
But if there was a real lab fire involving multiple chemicals and a firefighter accidentally sprayed 1 gallon per second of water onto a spilled pool of 4 L of concentrated sulfuric acid there would not be an issue. As far as the exothermicity of adding
water to solutions less than 50% there would not be enough heat generated to cause splattering issues.
Craig
Craig A. Merlic
Professor of Chemistry, UCLA Department of Chemistry and Biochemistry
Executive Director, UC Center for Laboratory Safety
Los Angeles, CA 90095-1569
Fascinating question.
H2SO4 is a strong acid which means it will immediately dissociate in water to yield H+ and HSO4-. It has a terrific affinity for water as we can see in common demonstrations such as sulfuric acid dehydration of sugar. Very exothermic,
which is why we always add acid to water and not vice-versa.
(HSO4- is a weak acid by definition as it does not dissociate completely to H+ and SO42-, so we only need concern ourselves with the first dissociation.)
In theory, an equimolar amount of water will fully consume the H2SO4 - there will be none left and we'd have H3O+ and HSO4-. However, while we say "completely dissociates" we are being a bit glib. The Ka1 is on the order of 10^3, so that
means the equilibrium lies strongly to the right, so this is an oversimplification. Regardless, the temperature spike on that would be nuts and vaporize some of the water unless this was done in a controlled fashion. So we need more water than equimolar. But
there's a lot more to this because we have to think about the enthalpy of mixing (heat of dilution). I could throw up a bunch of calculations but a more pragmatic approach is to ask if any diluted solutions of sulfuric acid are considered water reactive.
Commercial lead acid batteries have a 37% concentration. So I pulled up some random SDS's for lead acid batteries (of course, the authors may have no clue what they are taking about). Virtually all identified water as an incompatible or
noted that the material reacts violently with water. Therefore, this concentration is water-reactive.
Moving on to commercial 6 M solutions of sulfuric acid, a few SDS's mention avoid moisture, rather than listing it as incompatible, and some make no mention of any incompatibility. So, I would say this is the rule of thumb you're looking
for. Of course, this is for well-behaved systems like beakers and flasks, not tubs and drums and leaks in warehouses.
Can anyone help with this? I was asked to determine at what concentration would a sulfuric acid solution be considered no longer water reactive (or how dilute does it need to
be)? This is in the context of storage and firefighting/fire suppression systems (Building Code for Water Reactive Materials & NFPA 704 Annex F).
Maybe it is a simple calculation but my brain can't handle it today. Thanks!!
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