chem geek moments of the week

[baratron]

I love being a geek.

The other day I was walking down the road when I saw a van belonging to a cryogenic company. It had Hazchem plates saying it was transporting a non-hazardous non-flammable gas, and on the front was a sign in simple English saying "Delivering liquid carbon dioxide". I instantly went into chemistry geek mode, saying "well, I know carbon dioxide usually sublimes, so what sort of pressure must that carbon dioxide be under to be liquid?". And it's been bothering me for days, until just now when I had time to Google it and find out that the triple point of CO₂ is 5.2 atm. So, not actually all that much pressure.

Then lately I have been busy trying to solve a mystery relating to sulphuric acid. The Brønsted-Lowry theory of acids and bases refers to an acid as something which can donate an H+ ion (generally in chemistry referred to as "a proton" - although physicists will eeek at the thought of protons wandering around loose by themselves, which is further proof of why physicists and chemists can't talk to each other despite ~33% of their subjects being exactly the same material). Sulphuric acid is H₂SO₄, so it contains 2 hydrogens which can be lost as H+ (it's called a diprotic acid). So far, so good.

Now, the concentration of H+ in solution is useful for a lot of things, such as determining the strength of an acid or base, and measuring its pH. A strong acid is one which is fully split into ions in solution (compared to a weak acid where only a small proportion, e.g. 1/10000 of the molecules are split into ions). Therefore, the concentration of H+ (abbreviated as [H+]) for a strong acid is just the same as the concentration of the acid.

H₂SO₄ is a strong acid. Therefore, a lot of textbooks and web sites state that the concentration of H+ in sulphuric acid should be 2 x the concentration of the acid, because it's diprotic. I am wondering if this is actually true. Obviously, H₂SO₄ is a strong acid so the first dissociation is 100%:

H₂SO₄ (aq) --> H+ (aq) + HSO₄- (aq)

but then HSO₄- is a weak acid, and only partially dissociates:
HSO₄- (aq) --> H+ (aq) + SO₄2- (aq) (reversible)

The acid dissociation constant, K_a, for HSO₄- is, btw, 0.012.

So I think the [H+] for a 1 mol dm-3 solution of sulphuric acid should be 1.012 mol dm-3 : 1 mol dm-3 from the first dissociation, then 0.012 mol dm-3 from the second dissociation. This is a huge difference from the various books which claim that [H+] for a 1 mol dm-3 solution of sulphuric acid should be 2 mol dm-3. And I have no idea which of us are right.

I should write to the Royal Society of Chemistry about it for a definitive answer (says the person with MRSC after her name - if that doesn't qualify me to know what I'm talking about I don't know what does!).

Comments

[mhw.livejournal.com]

I knew that the critical pressure for CO2 couldn't be particularly high, since liquid Co2 is used as an industrial solvent, for example for dissolving caffeine out of coffee beans.

As to your question about [H3O+] (not a proton, but a hydronium ion) doesn't Le Chatelier's principle imply that the initial production of H3O+ from the first ionisation suppresses the secondary ionisation, due to HSO4- being a weak acid, so the [H3O+] won't be much different from that from the first ionisation alone.

I think. It's a long while since I did this stuff.

[baratron.livejournal.com]

I know that H+ is really H₃O+. The theory (and most textbooks/exam syllabuses) just refer to it as "a proton", though. It's one of those simple, satisfying and completely wrong definitions that are so ubiquitous in chemistry. Gotta love a science that's so imprecise at times.

Re Le Chatelier: It would, wouldn't it? So that's further evidence for me to wave at the people who are claiming that [H+] is 2 x [H₂SO₄]. Coolness!

I think a letter to the RSC's Education Division really is in order after this :) A lot of chemistry teachers need re-educating... ;)

[wuzzie.livejournal.com]

Should have asked me :(

Compressed liquid CO2 at ambient temperature will be somewhere around 55bar. Not especially high (my compressed air tanks operate up to 310bar) but it's still brown trowsers time when a burst disc blows :)

[hatter.livejournal.com]

You should have been in the kitchen when we were re-discussing it then 9as I'd also recently seen a similar van, and also been confused and surprised) - I'd assumed it would have been higher than your tank pressure, hence quite high, because I've not seen people rave about anodised or composite tank baffles to stop your pressurant slopping around as you run.


the hatter

[wuzzie.livejournal.com]

Ah, but the pressure will remain constant (for a given temperature) as long as there is liquid in the tank (assuming it's not been overfilled - "full" for a CO2 tank is only about 2/3 full of liquid.) It's just that it's not worth bothering with tank baffles when you're only carrying 20oz of CO2 around.

And I vaguely remember the conversation, but I thought it digressed rapidly onto the finer points of putting brains in the post...

[baratron.livejournal.com]

How do you put brains in the post? I googled for "how to post your brain", and then just "post your brain", and then "posting brains", and then "sending brains", and finally "mailing brains", and I still didn't find anything.

Enquiring minds need to know!

[ciphergoth.livejournal.com]

Ask these people? (http://www.brains4zombies.com/brains.html?brainType=OtherBrains&brainID=Helen-Louise)

[eub.livejournal.com]

A question for anyone who says it's 2 M protons would be "where does that leave the sulfate ions?" Kicking around loose? Nah.

AFAIR...

[bondagewoodelf.livejournal.com]

... molarity depends on the amount of known base you need for neutralizing the acid. The 1.012 mol dm-3 you give is the amount of H+ in solution at that concentration. If you start adding base, let's assume you added 1 mol/dm-3 of NaOH solution in water, you have 0.012 mol/dm-3 of H+ left plus the reaction products of the neutralization.

In essence you are confusing PH with acid molarity.

The 2 mol/dm-3 is correct. The 1.012 mol/dm-3 means the PH will approximately 0.99. If it had fully disassociated (if HSO4- would have been a stronger acid) PH would have been 0.70.

Re: AFAIR...

[baratron.livejournal.com]

Why would the molarity/concentration of an acid depend on the amount of base required to neutralise it? That wouldn't make any sense at all!

The concentration of an acid is the number of moles of acid in the solution (in a volume of 1 dm^3). The concentration of H+ is the number of moles of H+ in the solution, which may be more, less or equal to the concentration of the acid depending on whether it is a strong or weak acid.

I'm afraid it is you who is confusing pH with acid molarity :)

Re: AFAIR...

[bondagewoodelf.livejournal.com]

Why would the molarity/concentration of an acid depend on the amount of base required to neutralise it?

Because that's what molarity means: it's a measure for the amount of molecules or atoms in volume. It contains 2 mol/dm-3 of H+ in both associated -and- disassociated state.

Isn't PH a measure for the concentration of only the disassociated H+?

Anyway, 2 mol/dm-3 H₂SO₄^- (aq) is '2 mol/dm-3 sulfuric acid' which -has- 1.012 mol/dm-3 H⁺ (aq)

I guess you are right and the books are wrong when they say [H+] is 2 mol/dm-3

Re: AFAIR...

[baratron.livejournal.com]

You're missing a subtlety in my wording. Which is fair enough - your English is orders of magnitude better than my Dutch :) But my point is - Why would the concentration of an ACID be determined by the amount of BASE required to neutralise it? Surely the concentration of an acid (or any other chemical) depends ONLY on how much of the acid you have?

I think you're getting confused because you're half-remembering those horrible calculations where you have already mixed together an acid and a base and you're being asked to determine the concentration of each species and/or the pH of the mixture, but I'm not talking about that. I'm talking about the concentration of acid in a bottle of sulphuric acid that contains only sulphuric acid (and water) :)

Also, saying
It contains 2 mol/dm-3 of H+ in both associated -and- disassociated state.
is, er, rubbish (sorry), because the H+ only exists when the acid dissociates. Until then, all you can be certain of is that you have 2 mol dm^-3 (remember x^-y = x/y - mol per dm^-3 would be moles X dm^+3! (argh, I really need superscripts)) of HYDROGEN. 2 moles of H is not the same as 2 moles of H+ (aq)/H₃O+ ion. I mean, you could have 2 moles of H existing as part of OH-! Why not?

Re: AFAIR...

[bondagewoodelf.livejournal.com]

Why would the concentration of an ACID be determined by the amount of BASE required to neutralise it?

This come from me remembering that a '1 mol/dm³' afair means '1 molecular equivalent per cubic decimeter'. You do agree with me, that if you have a liter of sulfuric acid solution containing 1 mol of sulfuric acid, you need 2 liters of NaOH solution 1 mol/dm³ or 1 liter of NaOH solution 2 mol/dm³ to neutralize it?

It contains 2 mol/dm³ of H⁺ in both associated -and- disassociated state.

Maybe I should have said 'It contains acid in both associated -and- disassociated state so you need of the same amount of volume of 2 mol/dm³ NaOH solution to neutralize it'

Let's take it from this way: you have an acid solution, you have no idea which acid it is, but you want to determine it's concentration and all you have is 1 molar NaOH solution. So, you try a titration. After some fiddling to find out the correct indicator you realize you need one that changes color around PH 7. So you do the titration with this, and measure: for 10 ml of acid, I need 20 ml of this NaOH solution to neutralize the acid. Ergo, the original acid contains 2 mol/dm³. So, how much H₂SO₄ do you need to put into 1 dm³ to make this acid? 1 mol, right?

Please note I already agree with you that books claiming [H^{+/SUP>] for 1 molar sulfuric acid is 2 mol/dm3 are wrong. We are currently quibbling over my wording of my first reply ;-)}

[jinian.livejournal.com]

You are right; they are giant idiots if they're saying both protons dissociate. H₂SO₄ is a strong acid, but HSO₄- is not.

If molarity in British chemistry books somehow depends on what you need to neutralize an acid rather than what's actually there (which doesn't make sense to me because how would you have molarity of salt solutions then?), bondagewoodelf would be right, but that's certainly not like that in anything I've been taught. It could be a source of confusion, though. Are they saying outright that the molarity is 2 mol dm-3, or circumlocuting about base amount wanted to neutralize?

[baratron.livejournal.com]

Depends on the textbook - some are saying outright that [H+] is 2 mol dm^-3, others are simply saying that sulphuric acid is diprotic so needs 2 moles of alkali to be fully neutralised. The latter is true - or at least, always the way we do calculations. The former is... almost certainly NOT true - you could prove it fairly easily with a bottle of H₂SO₄ of known concentration and a pH meter.

I'm not sure that all the authors are giant idiots - some of them are quite good :) Actually, the person I emailed about it is someone who has written THE BEST chemistry textbook I've ever read (the kind of book I sat in bed and read cover-to-cover for pleasure rather than because I had to) - and it says quite a lot about this other book of his that I was pleased to find a possible error in it rather than saying "oh God, another stupid mistake"...