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CHRIS25
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Iodometry titrating for Fe3+ content
Object: to accurately assess amount of Fe2+ and Fe3+ in a solution of Iron Chloride.
In all the information I read about adding Potassium Iodide to the unknown analyte I read only that an excess should be added. This is vague for two
reasons:
1. How do you know when you have added excess to the Iron chloride solution?
2. What kind of molarity KI would generally be used?
and finally
3. What kind of molarity of sodium thiosulphate would generally be considered reasonable as a starter?
I need to have some idea about where to begin otherwise I could be wasting chemicals and repeating mistakes without knowing why. This sort of
information would be given by an instructor and in the absence of that I only have PDF's that give excellent and thorough explanations about what is
happening and why, but do not really help me to at least have a baseline from which to do this for the very first time. (and No I will not be
standardizing the thiosulphate solution with iodate, and will not be using some carbonate to stabilize the thiosulphate at this time).
Thankyou.
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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Quote: Originally posted by CHRIS25  | 1. How do you know when you have added excess to the Iron chloride solution?
2. What kind of molarity KI would generally be used?
and finally
3. What kind of molarity of sodium thiosulphate would generally be considered reasonable as a starter?
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A.1. You need to have some rough idea of the Fe3+ concentration in your sample. It should be somewhere between 0.05 M and 0.15 M, ideally. If
the concentration is much higher you will have to dilute the original sample to fall between the limits of 0.05 and 0.15 M, using a KNOWN dilution
factor. Then use 20 ml of this diluted solution as the sample to be titrated.
A.2. The exact molarity of the iodide is not important but assuming you'll be titrating an Fe3+ solution of about 0.1 M and that you add the iodide as
a 20 ml aliquot, it will have to be about 0.2 M to ensure there's an excess iodide, in accordance with 1.
A.3. The concentration of thiosulphate I would recommend here is 0.1 M (some may use 0.05 M or even 0.01 M). If the concentration of Fe3+ in 20 ml of
sample was about 0.1 M, you will need about 20 ml of titrant solution to reach end point.
But imagine the Fe3+ concentration was 0.5 M? The analysis would not work at all.
The approx. initial Fe3+ concentration can usually be worked out from the origin of the solution. Otherwise it becomes a case of trial and error.
[Edited on 15-5-2014 by blogfast25]
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CHRIS25
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A1 --- That is something I would have missed, but I presume this is because titrating iodine doe snot work at all in strongly acidic solns? But I
would never have guessed to have diluted it so much so Thankyou.
A2 --- All in all then Something that was completely absent from every doc I ever read, keep the Molarity of Fe soln low.
A3 --- Thankyou.
I know my present soln is around 3M to 4M and I know the amount of grams of iron, but I want to learn this method of titrating. Also since Iodide is
oxidized by the Fe3+ and is itself then reduced to Fe2+; it seems that any Fe2+ that happen to be in that soln
already would not be able to be calculated?
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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The Fe3+ concentration needs to be in the right area because otherwise your burette (25 ml) can't dispense enough titrant to react with all the
iodine. Since as you're using 0.1 M thiosulphate, the Fe3+ concentration in your sample should be in the same order of magnitude.
To dilute the 3 - 4 M original solution, take 5.0 ml and dilute it accurately to 250.0 ml (really you should use a suitable pipette and volumetric
flask for that task!). Or 10.0 ml to 500.0 ml. The dilution factor is then 50. Suppose your analysis came out as (e.g.) 0.0775 M on the diluted
sample, then original sample was 50 x 0.0775 M = 3.875 M.
Fe2+ cannot be detected with iodometry, unless you oxidise it to Fe3+, prior to the titration.
But for now stick to the Fe3+. Then we can work something out for the Fe2+... For instance titration of the Fe2+ with potassium permanganate.
[Edited on 15-5-2014 by blogfast25]
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CHRIS25
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Could someone Double check for me please
Method:
From a theoretical/suspected 3M soln of FeCl3 I extracted 0.1M. This was 4mLs.
Added 4mLs of FeCl3 to 116mLs water = Theoretical 0.012moles
Prepared KI 0.2M (0.004moles in 20mLs water)
Prepared S2O4 0.1M (0.0035moles in 35mLs water)
Prepared 0.75g corn flour in 100mLs boiling water, let cool, decanted the supernatant fluid
Titration Results:
When soln turned very light yellow corn flour was added and soln went deep blue. titration continued until deep blue disappeared at 17.5mLs.
However, although opaque at this stage the solution was not entirely colourless, continued until 26mLs had been added and then the soln turned truly
colourless with the clarity of water. I assume it is this mark that I should be calculating?
If so then I fall flat on my face with these calculations;
26mLs x 0.1M S2O4 / 4mLs FeCl3 = 0.65M
0.65M x 30 dilution factor = a very very concentrated FeCl3 soln of 19.5M???
Ok, I have really done my best to ensure complete accuracy, something is wrong and I simply can not see it.
(Assuming a 3M concentration in 125mLs that I have then the maximum can not be more than: 0.125 x 3 = 0.375moles of Fe which is close to the 0.5
moles that I originally started with; 19.5M gives me 2.4moles which is not possible).
[Edited on 23-5-2014 by CHRIS25]
[Edited on 23-5-2014 by CHRIS25]
[Edited on 23-5-2014 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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DraconicAcid
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| Quote: Originally posted by CHRIS25 | Method:
From a theoretical/suspected 3M soln of FeCl3 I extracted 0.1M. This was 4mLs.
Added 4mLs of FeCl3 to 116mLs water = Theoretical 0.012moles
Prepared KI 0.2M (0.004moles in 20mLs water)
Prepared S2O4 0.1M (0.0035moles in 35mLs water)
Prepared 0.75g corn flour in 100mLs boiling water, let cool, decanted the supernatant fluid
Titration Results:
When soln turned very light yellow corn flour was added and soln went deep blue. titration continued until deep blue disappeared at 17.5mLs.
However, although opaque at this stage the solution was not entirely colourless, continued until 26mLs had been added and then the soln turned truly
colourless with the clarity of water. I assume it is this mark that I should be calculating?
If so then I fall flat on my face with these calculations;
26mLs x 0.1M S2O4 / 4mLs FeCl3 = 0.65M
0.65M x 30 dilution factor = a very very concentrated FeCl3 soln of 19.5M???
Ok, I have really done my best to ensure complete accuracy, something is wrong and I simply can not see it.
(Assuming a 3M concentration in 125mLs that I have then the maximum can not be more than: 0.125 x 3 = 0.375moles of Fe which is close to the 0.5
moles that I originally started with; 19.5M gives me 2.4moles which is not possible).
[Edited on 23-5-2014 by CHRIS25]
[Edited on 23-5-2014 by CHRIS25]
[Edited on 23-5-2014 by CHRIS25] |
Ignore the dilution factor, since you titrated the whole thing anyway.
Okay, if the blue colour vanished at 17.5 mL, and your thiosulphate was 0.1 M, then you used 1.75 mmol thiosulphate. A 1:1 ratio means that you had
1.75 mmol Fe3+, which came from 4 mL solution, so that's 0.44 M Fe3+.
[Edited on 23-5-2014 by DraconicAcid]
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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CHRIS25
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So my maths, as is usual, came adrift. I get it thankyou, so waiting till the blue solution has gone and is clear, as is what I have read, is a
little misleading, one waits untill the blue disappears, regardless of the clarity. Dilution? That was stupid of me. Thankyou sincerely, Draconic.
Except this means my titration is off somewhere. 0.44M means only 0.05moles of Iron - this is way off. Or not all the ferrous has been oxidized?
But even with this factor considered and the colour of the solution all indicating oxidation of ferrous is more than 80% there.
Another titration tomorrow....
[Edited on 23-5-2014 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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| Quote: Originally posted by CHRIS25 | Method:
From a theoretical/suspected 3M soln of FeCl3 I extracted 0.1M. This was 4mLs.
Added 4mLs of FeCl3 to 116mLs water = Theoretical 0.012moles
|
Unfortunately both last sentences already don't make any sense. You can't 'extract 0.1M'.
Did you take 4 ml (four millilitres) of 3 M FeCl3 solution? If so what did you do with that?
Did you dilute that to 120 ml, i.e. '4mLs of FeCl3 added to 116mLs water'? Is that what you did?
General notational advice: ALWAYS, ALWAYS, ALWAYS separate a number from its unit of measurement by a space: e.g. 3 M, 5.6 g, 12 ml, 5.2 mol/L,
3.2 moles etc.
S2O4 0.1M ?? Did you mean 'Na<sub>2</sub>S2O3 0.1 M'? Preparing only 35 ml of that would be
folly: far too inaccurate.
[Edited on 23-5-2014 by blogfast25]
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CHRIS25
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====Did you take 4 ml (four millilitres) of 3 M FeCl3 solution? If so what did you do with that?
====Did you dilute that to 120 ml, i.e. '4mLs of FeCl3 added to 116mLs water'? Is that what you did?
Yes this is what I did.
==== ALWAYS, ALWAYS, ALWAYS ok point absorbed.
====S2O4 0.1M ?? Did you mean 'Na2S2O3 0.1 M'? Preparing only 35 ml of that would be folly: far too inaccurate.
woa I have no idea why I made such a bloody ignorant mistake (my actual notes have the correct formula); but yes, obviously I used sodium thiosulphate
pentahydrate
well two changes then to tomorrows titration; first I will make a 100 mL thiosulphate soln, and second, I will add the iodide, directly, as anhydrous
into the ferric soln, and not make a solution of potassium iodide and then add this.
[Edited on 24-5-2014 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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Chris:
From the dilution (dilution factor = 120 / 4 = 30), which amount did you then titrate?
Definitely not 4 ml, right? This where your error lies: in the calculation you need to use that number:
v<sub>thiosulphate</sub> x C<sub>thiosulphate</sub> = v<sub>diluted sample</sub> x C<sub>diluted
sample</sub> with v<sub>diluted sample</sub> the volume of diluted sample you actually subjected to titration.
Then multiply C<sub>diluted sample</sub> with the dilution factor to get the concentration of the original, undiluted sample.
What you need is a couple of decent (second hand will do) 250.0 ml Class A volumetric flasks and ditto 10.0 and 20.0 ml bulb pipettes and you'll soon
be up there with the pros!
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CHRIS25
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Now I am confused.
1. I do not know the concentration of the diluted sample since I do not know the number of moles in the 4 mL sample taken. It was 4 mLs + 116 mLs
water. And since I titrated into this whole solution, (not a part of it) I thought that what Draconic said was right - I do not need to bother with
concentration. In fact when I titrate sodium hydroxide into a diluted copper chloride solution I never calculate the dilution factor into it, and
luckily I saw many demonstrations that did not calculate concentration in this instance (copper chloride etching process).
2. I do not know what you mean by the word "with" in (with v diluted sample); which mathematical notation is the 'with' referring to?
3. I will deal with extra lab stuff later, but for now I have to use syringes for complete accuracy, (they are so much cheaper than bulb pipettes),
and since they are from a veterinary source I can hardly doubt their accuracy for my needs at this moment in time.
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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Chris:
Normally one would never titrate the entire dilution but only a known part of it, typically 20.0 ml. Otherwise you risk having to use far, far too
much titrant solution to be practical.
Anyway, you titrated the whole 120 ml (containing the entire 4 ml of sample), so lets work with that.
You used 26 ml of 0.1 M Na2S2O3, so that you calculate the number of mol thiosulphate used being (0.026 L x 0.1 M) = 0.0026 mol.
But already there's a snake in the grass.
The titration reaction is:
S2O3(2-) + I2 === > 1/2 S4O6(2-) + 2 I<sup>-</sup>
And the original reaction:
Fe3+ + I<sup>-</sup> === > Fe2+ + 1/2 I2
In other words, each mol of thiosulphate titrates for TWO (2) mol of Fe3+ !!
So you need to double that 0.0026 mol of thiosulphate to 0.0052 mol of Fe<sup>3+</sup>. Now divide that by the volume in which they were
contained: 4 ml or 0.004 L: 0.0052 mol / 0.004 L = 1.3 M Fe<sup>3+</sup>
Still quite a distance from the suspected 3 - 4 M. I'll later look at the amount of KI added, to see if it is correct. How much of the 0.2 M KI
solution did you add to your titrated sample?
Two general remarks: the doubling you overlooked is why experienced titrators use NORMALITY instead of MOLARITY.
Secondly, using syringes and such like is fine but difficult. Experienced titrators use a simple protocol which involves Normal solutions (of correct
strength), a 10 or 25 ml burette, volumetric flasks and calibrated pipettes, because small errors soon mount to quite a lot and it can be difficult
finding out what went wrong. QED.
Take this for example: "Prepared S2O4 0.1M (0.0035moles in 35mLs water)". That's 0.86863 g of sodium thiosulphate pentahydrate, difficult to
accurately weigh without using 0.1 mg scales.
I just prepared a 0.1 N Na2S2O3 solution, 250.0 ml of it. That's about 3.1 g of sodium thiosulphate pentahydrate, which I can accurately enough weigh
with my 1 mg scales (measuring error is about 0.3 %)
[Edited on 24-5-2014 by blogfast25]
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CHRIS25
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I stared at the balanced equation for this reaction and saw that one mol of FeCl3 resulted in half the amount of iodine but I assumed
something else wrongly in my thinking.
Yes you are right, I weighed out 0.9g of thiosulphate because my scales do not go above one decimal place. (cheap ebay), anyway I realise now that
measuring in such small quantities magnifies errors more than measuring in large quantities. (a case of economics at home); but I will do the
titration again with larger amounts and with probably 10g of ferric solution and use 0.15 M of thiosulphate extracted from 100 mLs I suppose.
I poured all 20 mLs of the 0.2 M KI into the ferric solution, this I measured out at 0.66g but put in 0.7g and then pinched a bit out of the container
in an attempt to lower it closer to 0.6 while still on the scales, (scales again yes I know new scales would be better). The mistakes you have
highlighted though - thanks.
[Edited on 24-5-2014 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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Ouch: 1 decimal place! Too little. The only way to use that quasi accurately is by doing everything on a 10 x scale factor. E.g. 9 g of thio in 10 x
the amount of water. A totally false economy.
2 digits isn't great but it's workable, the ones I sell cost just under £10. Can't go wrong with that. Better: 0.001 g (about £50, I think).
Not sure what you mean by: "[...] and use 0.15 M of thiosulphate extracted from 100 mLs I suppose."
[Edited on 24-5-2014 by blogfast25]
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CHRIS25
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Actually I just measured 0.15 M ready for 200 mLs of water; extracted? I meant merely that I am making a beaker with this amount ready to use, 7.44g
in 200 mLs = 0.03 m. 10 pounds? Mmm.
[Edited on 24-5-2014 by CHRIS25]
on a side note yes it is nice to have good tools for the job, but I suppose for me at the moment the learning and doing is so much more important,
even if accuracy is temporarily sacrificed - the practise with what I have is important.
[Edited on 24-5-2014 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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0.15 M, so you mean 0.15 mol thio per litre. Why are you increasing it from 0.1 to 0.15 M? It would be slightly more to weigh, so a little more
accurate. Is that the purpose?
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CHRIS25
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| Quote: Originally posted by blogfast25 | | 0.15 M, so you mean 0.15 mol thio per litre. Why are you increasing it from 0.1 to 0.15 M? It would be slightly more to weigh, so a little more
accurate. Is that the purpose? |
Yes I thought a little more leads to just a little more accuracy?
0.15 M x 0.2L = 0.03 mol. 0.03 mol x 248 g/mol = 7.44g
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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| Quote: Originally posted by CHRIS25 | | Quote: Originally posted by blogfast25 | | 0.15 M, so you mean 0.15 mol thio per litre. Why are you increasing it from 0.1 to 0.15 M? It would be slightly more to weigh, so a little more
accurate. Is that the purpose? |
Yes I thought a little more leads to just a little more accuracy?
0.15 M x 0.2L = 0.03 mol. 0.03 mol x 248 g/mol = 7.44g |
Yes, it's a little better: 0.1 / 7.4 x 100 % = 1.8 % relative measuring error. Much better than 0.1 / 0.9 x 100 % > 10 %.
Spare a thought for the men and women who determined atomic masses to 5 decimal places!
[Edited on 24-5-2014 by blogfast25]
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CHRIS25
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Relative measuring error is something totally new to me. Is it a standard that chemists use? And the way you demonstrated above is the maths for
this?
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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| Quote: Originally posted by CHRIS25 | | Relative measuring error is something totally new to me. Is it a standard that chemists use? And the way you demonstrated above is the maths for
this? |
The more I do error analysis, the more I realise just how damn hard it is to measure ANYTHING with known and sufficient accuracy.
The relative measuring error is just a basic comparative tool.
For more details on error analysis, consult some basic texts on the Tinkerwebs. It ranges from fairly basic algebra to full blown 'Monte Carlo'
statistical analysis.
[Edited on 24-5-2014 by blogfast25]
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CHRIS25
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ok, thanks , but 0.1 / 7.4 = 0.0135, 0.0135 x 100 = 1.35?
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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Yup. Typo on my calculator. Another rocket that would go after its creator! ;-)
[Edited on 24-5-2014 by blogfast25]
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aga
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How was that done ?
I have also searched for exactly HOW chemical forumulae and compound composition were deduced in the 1800s and have drawn a blank.
Any pointers gratefully received.
[Edited on 24-5-2014 by aga]
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blogfast25
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| Quote: Originally posted by aga |
How was that done ?
I have also searched for exactly HOW chemical forumulae and compound composition was deduced in the 1800s and have drawn a blank.
Any pointers gratefully received.
[Edited on 24-5-2014 by aga] |
When French chemist Louis-Nicolas Vauquelin in 1798 discovered beryllium (in the mineral Beryl), he determined the atomic mass to be about 9. He was
right but today we know this with astonishing accuracy: 9.0121831 g/mol. How did we get from simply "9" to 9.0121831? That is a very long story of
technological evolution and ever refining of available analytical and measurement techniques and not something that can be easily summarised here.
Similarly the story of stoichiometry, that goes back almost to the ancient Greeks who were the first to propose the idea of an atom, after which the
idea went dormant for several centuries.
Gradually and over several hundreds of years, competing 'philosophies' (as they were then called) vied to provide the best explanations for the
bewildering amount of compounds and materials we see in the world. Ultimately the so-called 'atomists' won and the notion that chemical elements
combine into compounds in fixed ratios became established.
That part of the story, intertwined with the first is a picture so large that you can't even fit it onto a single canvass (so to speak).
The history of science is almost as compelling as the science itself. Today there's no great shortage of great books that popularise the history of
physics and chemistry. I'll see if I can rustle up some great titles.
In terms of popular science history shows, I like Jim al Khalili's 'Chemistry: a Volatile History' a lot. Regularly repeated on Beeb 4.
[Edited on 24-5-2014 by blogfast25]
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aga
Forum Drunkard
Posts: 7030
Registered: 25-3-2014
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I have just searched for Louis-Nicolas Vauquelin and again learnt nothing as to HOW the chemical compositions or atomic weights were deduced.
Am i being thick ?
Can everyone else pick up a lump of charcoal and say :-
'Hmm. Thats Carbon that is. Atomic weight's about twelve i reckon. Dunt feel like 14 isotope to me, more like 12' ?
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