I haven't really touched or looked into anything regarding coordination complexes, but would like to see if I can get in to it. Does anyone know some
simple coordination complexes that would be easy to make as an introduction to them?Metacelsus - 11-3-2015 at 10:33
Tetraamminecopper(II) complexes are fairly easy to make.Sulaiman - 11-3-2015 at 11:18
then I look for publications etc. for whatever particular aspect interests me.
Very often whilst researching via google I find references to this Science Madness forum!
[Edited on 11-3-2015 by Sulaiman]DFliyerz - 11-3-2015 at 14:31
I also have an additional question, which probably seems sort of weird, but I can't help but wonder it: do all coordination complexes act like ionic
compounds where they can be crystallized and exist outside of solution, or can only certain complexes do this?DraconicAcid - 11-3-2015 at 15:27
The stable ones exist outside of solution. Neutral coordination complexes (such as Ni(DMG)2 or Cr(acac)3) often act as molecular compounds.Magpie - 11-3-2015 at 19:05
Here's two I like:
1. Fehling's test reagent for aldehydes.
Part I: dissolve 34.64g of CuSO4.5H2O in 350-400 ml water. Dilute to 500 mL.
Part II: dissolve 173g Rochelle salt and 65g NaOH in ~350 ml of water. Dilute to 500 ml.
The test: To 3 ml of part I slowly add part II. Initially a light-blue precipitate of copper hydroxide forms. But eventually this dissolves and a
dark blue complex of the tartrate ion forms upon shaking.
2. Tollen's test reagent for aldehydes
30g of AgNO3 is dissolved in 500 ml of water. A solution of NH4OH is slowly added until the ppt of silver Ag2O that first forms is barely
redissolved.
I believe that NH3 combines with the Ag+ to form a colorless complex.
Source: lab manual of Brewster et al Pasrules - 12-3-2015 at 01:32
I made Aquabis(glycinato)copper(II) last week from Copper(II) sulfate pentahydrate. Lovely Blue
And today i made Tris(en)cobalt(III) from cobalt(II) chloride hexahydrate. Lovely Orange, currently in the dessicator.
Real shame i can't take photos in the lab.bismuthate - 12-3-2015 at 03:39
Cobalt is really great for complexes. You can just add HCl to a Co2+ solution and get the chloro complex. Also if you have and thiocyanate that forms
complexes very easily with metals. If you have many transition metals, some of them (like titanium and chromium) form complexes with H2O2.woelen - 12-3-2015 at 03:57
Try to obtain ethylene diamine (en). With that chemical you can make many amazing complexes. If you want to crystallize complexes, then it also is
nice to have a highly soluble perchlorate (best is NaClO4, second best is NH4ClO4, KClO4 is not useful). In this way I have made solid pure (en)
complexes of copper(II), cobalt(III), nickel(II), zinc(II), with perchlorate as counter ion. Many complex perchlorate salts are sparingly soluble and
can easily be crystallized.Amos - 12-3-2015 at 04:21
Tetraammine copper(II) sulfate and hexaammine nickel(II) chloride are both easy to make and can be isolated in solid form and kept.
Tetraammine copper(II) sulfate can be produced by dissolving copper(II) sulfate in the minimum amount of water necessary, and adding aqueous
ammonia(as strong as you can find, the idea is to keep the amount of water low). A precipitate will form this way, but continue adding the ammonia
with stirring and the precipitate will dissolve, forming a brilliant royal blue complex. Once no more solid remains, slowly pour in acetone while
stirring; the complex will come out of solution. The precipitate forms large particles and can easily be filtered out, and if you're quick enough you
can dry it using a device that provides a stream of hot air, like a hairdryer.
Hexaammine nickel(II) chloride is produced very similarly to the above, only nickel(II) chloride is used in the place of copper(II) sulfate. The
addition of ammonia to the chloride is actually rather exothermic, so I recommend that the reagents are at least kept cool. Do remember that nickel
salts are toxic and carcinogenic! Use gloves if you do this!Metacelsus - 12-3-2015 at 14:32
About drying the tetraammine copper(ii) sulfate: too much heat will cause it to lose ammonia gas and become copper(ii) sulfate. I don't recommend
heating it with a hair dryer.DraconicAcid - 12-3-2015 at 16:10
About drying the tetraammine copper(ii) sulfate: too much heat will cause it to lose ammonia gas and become copper(ii) sulfate. I don't recommend
heating it with a hair dryer.
The nickel compound loses ammonia even more readily.woelen - 12-3-2015 at 23:50
I myself made the tetrammine complex of copper sulfate.
I did the following:
Dissolve as much as possible of copper sulfate in water. Make the solution warm to get even more in solution.
Drip in ammonia (25%). First you get a blue precipitate, but this redissolves, giving a really dark blue solution. Add a little excess ammonia, such
that the solution has a clearly noticeable smell of ammonia.
Add ethanol (appr. same volume as the volume of the aqueous solution). Denatured ethanol is OK (e.g. spiritus ketonatus), but it must be good stuff,
not the dyed stuff, which is used for cleaning windows.
Set aside for a few hours. Long needle-like crystals separate.
Put in the fridge for another few hours to get more crystals.
Decant the liquid from the crystals. Put the crystal mass on a folded coffee filter, which itself in turn is put on a pile of paper tissues or pieces
of toilet paper. In this way you already get fairly dry crystals.
Rinse the crystals with acetone (or if you have that, with diethyl ether) and again put the crystal mass on a piece of coffee filter with paper tissue
underneath.
Transfer to a dry filter and allow to dry for half an hour or so. Then store in tightly sealed container.
The result is as follows:
