Could someone kindly check this Titimetry figure please, and probably the method as well.
It's my first time ever, so am unsure that i have the method correct, or that the result is sane.
Today i made some ammonia solution, and wished to see what the concentration was. It certainly makes my eyes water, so it's certainly more
concentrated than OTC stuff.
The only item in my collection with a Molar concentration figure on the label is Nitric acid, 6.8 [M], so i used that as the Titrator, given that NH3
solution is basic.
Knowing little, i am assuming that 1 Mol Nitric Acid neutralises 1 Mol Ammonia.
To begin with i diluted the NH3 with some water, and the same with the HNO3, then added some pH Indicator to the NH3 solution and gradually added the
Nitric acid solution to see if it fizzed, popped, heated etc, and crucially to see if the colour changed.
The colour did change, with no fizzing or popping or heating, so i had a rough idea of the volumes/dilutions required to do the test, and what the
reaction would do.
1ml of the 6.8 [M] HNO3(aq) was diluted with distilled water to a total volume of 50ml.
1ml of the NH3(aq) and some red cabbage water pH indicator were diluted to a total volume of 40ml.
At this point the Titrand solution + indicator were dark green.
The HNO3 solution was gradually added to the NH3(aq) & Indicator solution using a 5ml syringe.
At about 35ml, the colour began to change, and patches of purple/violet could be seen. Oddly they occupied areas in the center of the vessel, and
wouldn't move despite swirling. Guess i need to study fluid dynamics to work that one out.
The colour matched (best my eyes could see) the Control indicator sample (violet) in a test tube of pure distilled water, when 45ml of the nitric acid
solution was added.
My calculations make this about 4.9 [M] ammonia solution, but i stand to be (repeatedly) corrected.
If that's right, is that concentrated enough for the tetraammine (II) copper complex reaction ?blogfast25 - 2-4-2014 at 11:41
Aga:
Use c<sub>1</sub> x V<sub>1</sub> = c<sub>2</sub> x V<sub>2</sub>
For the dilution of the nitric acid: 6.8 mol/1 L x 0.001 L = c<sub>2</sub> x 0.05 L
So c<sub>2</sub> = 0.136 M (mol/L)
I’ll assume your end point of 35 ml was correct. In that case you used 0.136 mol/L x 0.035 L = 0.00476 mol HNO3 to neutralise the NH3:
HNO3 + NH3 === > NH4NO3 + H2O
… so one mol of HNO3 corresponds (is ‘equivalent’ as we say) to one mol of NH3.
The 0.00476 mol of NH3 were contained in 1 ml = 0.001 L, so the concentration is 0.00476 mol NH3 / 0.001 L = 4.76 M NH3.
Blimey, not bad at all for a first attempt!
There is no real need to compare the colour to that of a pure water solution (with indicator): in titrometry with indicators you look for an overall
colour change. Comparison to some standard colour can be quite misleading. It's easier also to use synthetic indicators: they show sharper
end points. For NH3/HNO3 methyl orange is one of the recommended ones.
Yes, it’s probably strong enough for the copper ammine complex but make sure you start with a weak CuSO4 solution, for instance 0.1 M or less. Add
to it about the same amount (volume) of your ammonia. Stir well, perhaps stand overnight.
Two reactions have to take place: firstly formation of Cu(OH)2 and then dissolution of the Cu(OH)2 in excess NH3, to form the complex.
[Edited on 2-4-2014 by blogfast25]elementcollector1 - 2-4-2014 at 11:51
Our experiment with the same took roughly 15 minutes - no idea on the concentration of the ammonia, but I seem to remember it being household. This is
most easily evidenced by a striking blue-purple, transparent solution, to which ethanol was added to force precipitation of a dark blue-purple sludge,
which was vacuum-filtered to yield a cake of blue powder. Personally, I think the solution was much more beautiful than the powder, which is why I put
it on 'Pretty Pictures' (somewhere...).aga - 2-4-2014 at 11:55
Many thanks blogfast25.
There is a warm glow occurring ...
Some chemicals, a syringe, a red cabbage, some jam jars and a result !
I noticed the violet patches at 35ml, but the total change was at 45ml.
Was 35ml the 'end point' rather than at 45ml ?DraconicAcid - 2-4-2014 at 12:02
Some chemicals, a syringe, a red cabbage, some jam jars and a result !
I noticed the violet patches at 35ml, but the total change was at 45ml.
Was 35ml the 'end point' rather than at 45ml ?
No- if you have patches of one colour, that means the solution is not properly stirred. The entire solution has to change colour before you can
declare it the endpoint.aga - 2-4-2014 at 12:05
Note to self : Stir
Subnote: Draconic Stirring !
[Edited on 2-4-2014 by aga]aga - 2-4-2014 at 12:36
... make sure you start with a weak CuSO4 solution, for instance 0.1 M or less.
Ah.
So CuSO4(aq) titrimetry tomorrow then.
This time stirring will happen.
Many thanks peeps.blogfast25 - 3-4-2014 at 04:30
Yes, the entire liquid needs to have changed colour and your end point was probably close to 45 ml, rather than 35. Stirring is what keeps things
homogeneous, and so nearly always necessary (where possible).
When you do the copper ammine complex, hold back two volumes of the CuSO4 and to one of them add some table salt and shake up. Observe. aga - 4-4-2014 at 05:29
Doh ! I didn't make enough ammonia solution so the tetraammine thing didn't work.
Undeterred i just kept adding OTC 3% ammonia solution to the CuSO4(aq) and eventually it all went nice and Deep Purple (but with no smoke on the
water).
Inherrently suspicious i added some NaCl (as suggested) to 10ml of the tetraammine complex in a test tube and shook it up to see if any black holes or
goblins appeared. They didn't.
I re-read the suggestion, and added NaCl(s) to CuSO4(aq).
The solution went green with no fizzing or heating.
I guess it went to copper chloride and sodium sulphate ?
There's a solid at the bottom of the tube, but i can't tell if it is a precipitate, or just some remaining NaCl, so i'll try it again with a solution
of NaCl.blogfast25 - 4-4-2014 at 09:33
The green is caused by yet another complex of Cu(II): CuCl<sub>4</sub><sup>2-</sup>, aka tetrachlorocuprate, because you've
added NaCl.
So far you have seen three complexes of copper: the one above, the tetrammine (deep blue) and the so-called hexa aqua complex. The latter is what
gives normal dissolved CuSO<sub>4</sub> its sea blue colour:
[Cu(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup>. But when you add chloride anions to that, the latter displace the
H<sub>2</sub>O molecules.
That doesn't happen when you add chloride to the tetrammine complex because the latter is 'stronger' (more stable) than the chloro complex. In other
words for the reaction 'ammine complex === > chloro complex' the Delta G > 0 (cannot proceed).
The big, middle block of the periodic table (what we call the 'D-block') is full of coloured compounds, often due to similar complexes. Nickel and
Chromium for instance form colourful complexes with water, ammonia and chloride. Other complexing agents also exist.
