Sciencemadness Discussion Board - Normality and molarity

Normality and molarity

Little_Ghost_again - 17-11-2014 at 02:39

Hi
While reading a procedure I came across something thats got me wondering.
It spoke of a synth with Nitric acid and described a 1M 1N solution of nitric acid.

I thought a 1N solution of Nitric acid would be 0.5 Mols as it has 2H+ or does it?

Ifa solution is described as 1N then surely you would need 0.5M of 70% solution.
BTW the conc was given as 70% Nitric Acid
Anyone throw some light on this, I dont have a link as it was a stumble page and I only thought about it after

Chemosynthesis - 17-11-2014 at 02:56

Monoprotic.
HNO3.

Little_Ghost_again - 17-11-2014 at 03:22

I messed up I meant H2SO4, but was doing something this morning with Nitric acid, so replace all reference to Nitric with sulphuric! sorry about that

Chemosynthesis - 17-11-2014 at 03:27

Generally speaking, 1M would then by 2N, yes... but that is only when assuming titration of hydroxide or similar. I suppose it may be able to refer to 1M sulfuric acid as 1N if only the first protonation were of a low enough pKa to dissociate under specific conditions, but I haven't seen this.

blogfast25 - 17-11-2014 at 05:54

The concept of normality ('equivalents per litre') is also used in redox titrations. For instance 1 mol of dichromate reduced to 1 mol of Cr(III) in acid conditions absorbs 6 mol electrons. A 1 M solution would thus be 6 N. In reality you would normally titrate with a 0.1 N solution or 0.1/6 M. Or a 0.05 N, i.e. 0.05/6 M solution.

They don't seem to teach normality all that much nowadays, which is a shame because in many reactive situations it is easier than molarity: 1 equivalent of whatever always reacts with 1 equivalent of whatever else.

[Edited on 17-11-2014 by blogfast25]

Little_Ghost_again - 17-11-2014 at 06:25

Hi
actually yes I didnt mention the context!! Sorry having one of those days! It was to do with titration which is why I was confused.
Thanks again
LG

Brain&Force - 17-11-2014 at 09:57

Quote: Originally posted by blogfast25

The concept of normality ('equivalents per litre') is also used in redox titrations. For instance 1 mol of dichromate reduced to 1 mol of Cr(III) in acid conditions absorbs 6 mol electrons. A 1 M solution would thus be 6 N. In reality you would normally titrate with a 0.1 N solution or 0.1/6 M. Or a 0.05 N, i.e. 0.05/6 M solution.

They don't seem to teach normality all that much nowadays, which is a shame because in many reactive situations it is easier than molarity: 1 equivalent of whatever always reacts with 1 equivalent of whatever else.

[Edited on 17-11-2014 by blogfast25]

I am so glad I read this thread for this very reason...

DraconicAcid - 17-11-2014 at 10:05

Meh. I see "normality" as an extra layer of confusion added for the benefit of people who don't understand stoichiometry.

blogfast25 - 17-11-2014 at 10:50

Quote: Originally posted by DraconicAcid

Meh. I see "normality" as an extra layer of confusion added for the benefit of people who don't understand stoichiometry.

No, no: you can't work normalities if you don't understand stoichiometry but in that case you can't do it with molarities either!

Where I see it as useful for instance is where a 'chain' of redox reactions is used for titration: e.g. oxidise ethanol with excess dichromate, oxidise iodide to iodine with left over dichromate, titrate iodine with thiosulphate.

Determine valences for each species (half reactions) and prepare equinormal solutions of all, e.g. 0.1 N for all solutions. Now you know 1 ml of one always corresponds to 1 ml of the others.

I know some find it confusing but they remind me of those Brits who still can't convert pounds to kilogram!

Little_Ghost_again - 17-11-2014 at 10:55

Quote: Originally posted by blogfast25

Quote: Originally posted by DraconicAcid

Meh. I see "normality" as an extra layer of confusion added for the benefit of people who don't understand stoichiometry.

Mols I get...................I am also a Brit so maybe thats why?

DraconicAcid - 17-11-2014 at 10:56

Quote: Originally posted by blogfast25

I know some find it confusing but they remind me of those Brits who still can't convert pounds to kilogram!

I've met Brits who couldn't even convert stones to pounds. (They were selling tickets to use these human-sized hamster-ball things, but had a sign up saying that you had to be under a certain number of stone to use one. They couldn't tell me how many kg that was, or how many pounds. It took me five minutes to get them to decide if I was light enough to go in one.)

unionised - 17-11-2014 at 11:57

Quote: Originally posted by blogfast25

They don't seem to teach normality all that much nowadays, which is a shame because in many reactive situations it is easier than molarity: 1 equivalent of whatever always reacts with 1 equivalent of whatever else.

[Edited on 17-11-2014 by blogfast25]

Except when it doesn't.
For example, a solution of KMnO4 that is 1M is 5N if you are doing the titration in acid and 3N if you are doing it in base.

Good luck with things like peroxyacetic acid and formic acid.
They can be titrated as acids or as an oxidant and reductant respectively.

Magpie - 17-11-2014 at 12:38

Quote: Originally posted by blogfast25

No, no: you can't work normalities if you don't understand stoichiometry but in that case you can't do it with molarities either!...

I agree. I always use normality when doing titration calculations.

Quote: Originally posted by blogfast25

I know some find it confusing but they remind me of those Brits who still can't convert pounds to kilogram!

Yes, in practical every day usage 454g = 1 lb.

But technically, one unit is mass and the other is force!

unionised - 17-11-2014 at 13:23

Well, in the foot pound second system the unit of force is the poundal so I guess you think that the kilogram is a unit of force.
Or is this the unit of force you meant
http://en.wikipedia.org/wiki/Slug_(mass)

aga - 17-11-2014 at 13:25

Aside from the mathematical utility of Normality, it does make you think about the specific reaction more carefully.

Personally i found the use of Normality in a recent series of titrations a refreshing change, as there was no Normality around here before.

Etaoin Shrdlu - 17-11-2014 at 13:44

Quote: Originally posted by Magpie

Quote: Originally posted by blogfast25

No, no: you can't work normalities if you don't understand stoichiometry but in that case you can't do it with molarities either!...

I agree. I always use normality when doing titration calculations.

Quote: Originally posted by blogfast25

I know some find it confusing but they remind me of those Brits who still can't convert pounds to kilogram!

Yes, in practical every day usage 454g = 1 lb.

But technically, one unit is mass and the other is force!

Except, in the US at least, the pound is also a legal unit of mass!

I got into a lighthearted argument with a lawyer over weight vs mass and was flabbergasted to learn this. Not that it doesn't make sense on some level, it's just in the sciences I've always seen the pound as a unit of force and I've had it drilled into my head forever (with him it was the exact opposite).

I think this is true in some international standard as well.

Quote: Originally posted by unionised

Well, in the foot pound second system the unit of force is the poundal so I guess you think that the kilogram is a unit of force.
Or is this the unit of force you meant
http://en.wikipedia.org/wiki/Slug_(mass)

This one. http://en.wikipedia.org/wiki/Pound_(force)

[Edited on 11-17-2014 by Etaoin Shrdlu]

Little_Ghost_again - 17-11-2014 at 14:04

I am sorry I asked now

blogfast25 - 17-11-2014 at 14:16

Quote: Originally posted by unionised

That's an old objection that's not exclusive to normality.

The specific reaction determines the valence used to calculate the equivalent mass of the species. When using molarities you need to take that into account too. Someone who doesn't understand the reductions of permanganate in acid or alkaline conditions is likely to get it wrong, molarities or normalities no matter. It's not a problem inherent to the idea of using equivalents instead of mol.

If you titrate H3PO4 with NaOH, which end point will you choose? Same problem, same easy solution: think before you begin, lest you end up getting it wrong by a factor two or three.

The same is true for oxalic acid (as acid or reducing agent) and the other examples you bring up: If you don't even know what reaction you're using them in you shouldn't be in a laboratory but in an embroidery course.

Quote: Originally posted by Little_Ghost_again

I am sorry I asked now

It's always been of these useless points of debate ('N v. M'): you gave some here something to bitch about, so don't worry!

[Edited on 17-11-2014 by blogfast25]

aga - 17-11-2014 at 14:59

Surely Molarity and Normality are both useful Tools to have in one's toolbox, so any argument about which is the Better Tool is utterly pointless ?

Just use the one that suits your situation best.

('situation' including which one you like best).

Little_Ghost_again - 17-11-2014 at 16:20

I confess when I first saw 1N on here I thought is was a typo

at least slowly I am gaining knowledge. Every day I learn something new is a happy day