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: 2004-10-22 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"...