chem geek moments of the week
Oct. 19th, 2004 01:34 am![[personal profile]](https://www.dreamwidth.org/img/silk/identity/user.png){kind=small}
baratronI 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, Ka, 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!).
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, Ka, 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!).
![[identity profile]](https://www.dreamwidth.org/img/silk/identity/openid.png){kind=small}
no subject
Date: 2004-10-19 05:50 pm (UTC)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?),
no subject
Date: 2004-10-22 05:59 pm (UTC)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"...