ScienceBum
Harmless
Posts: 22
Registered: 15-9-2014
Location: Florida/California/Arizona
Member Is Offline
Mood: decoherenced
|
|
Question regarding Hydration and Ions in Solution
I recently started studying chemistry again after finding my old college textbook, and in it is described ionic compounds in solution and hydrated
ionic compounds. I don't quite understand the difference, and without going into complex stuff like lattice energies and whatnot that I've been trying
to look up online, if someone could enlighten me on this I would very much appreciate it.
From my understanding, Ionic Compounds in Solution are completely dissociated into their respective ions due to the polar bonding
nature of water molecules - the charges are strong enough to completely separate the ions(so dissolving salt in water would get you Sodium and
Chlorine ions bonded to the water molecules as opposed to solid Sodium Chloride). Presumably, after evaporating the water the sodium and chlorine
would react together again to leave an equal amount of salt that was originally dissolved in the water.
The small section that I read regarding Hydrated Ionic Compounds said something along the lines of the ionic compounds having a
specific number of individual water molecules bonded to them, even being able to have "half" of a water molecule. The example in my book showed
pictures for comparison of Cobalt(II) Chloride hexahydrate and it's anyhydrous form, with the anhydrous compound being a different color.
Does hydrating an ionic compound change just it's physical properties, or both chemical and physical? If so, are there general "types" of properties
hydrating an ion will give it or is that dependent on various factors?
If anyone can help, I would greatly appreciate it. Thanks all
- cap'n rum the science bum
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
Hydration can change the chemical properties of a salt. I've worked with terbium chloride hexahydrate, which hydrolyzes to terbium(III) oxide and
hydrogen chloride when heated. Anhydrous terbium(III) chloride doesn't do this.
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
ziqquratu
Hazard to Others
Posts: 385
Registered: 15-11-2002
Member Is Offline
Mood: No Mood
|
|
Most definitely hydration can have a big effect on the properties of a compound. The term you're looking for to help your searches is "crystal field
theory". This is the theory which describes the effects of ligand binding on the properties of a metal complex. Water is a pretty decent ligand, and
of course coordinates to the (positive) metal ion via the (partially negative) oxygen atom.
It's been a while since I've had to remember all this, but essentially the binding of a ligand alters the energy levels of the various d and f
orbitals in the metal ion, which in turn can have big effects on the properties of the metal ion. The colour can change from anhydrous to hydrated,
and cobalt chloride is a good example, in fact - it is (was?) used as an indicator in silica gel used in, for example, dessicators, since the colour
change tells you when the gel has absorbed as much water as it can. Copper sulfate is also a good example - the common blue crystals are the
pentahydrate, whilst the anhydrous compound is nearly colourless. The colour change arises because of the change in orbital energy levels, which means
different wavelengths of light will be required to excite the electrons in those orbtials.
Ligands can affect the magnetic properties of metal complexes, too - because of the change in the energy levels, the electrons will shuffle around
into the lowest energy configurations. If this results in complexes with all electrons paired, the compound will be diamagnetic; with a different
ligand, though, there might be unpaired electrons, and thus a paramagnetic compound.
Ligands definitely affect reactivity, too - which is the basis of the entire field of coordination chemistry. A simple example, many metal complexes
are strong Lewis acids in their anhydrous forms, but not when hydrated. There are, however, plenty of more extreme examples - one I've worked with is
ruthenium-catalysed hydrogenation. When you use only phosphine ligands, ruthenium complexes can be great catalysts for the hydrogenation of alkenes,
but suck for most ketones. Adding an amine ligand, though, completely reverses that pattern - those complexes are great for ketones and barely react
with alkenes at all!
As I said, my coordination chemistry is quite rusty these days, but I think I've got the broad strokes correct (hopefully someone will step in if I've
messed something up!). As I said, if you're interested in more detail, have a search around and you should find plenty of good material to read.
|
|
|
ScienceBum
Harmless
Posts: 22
Registered: 15-9-2014
Location: Florida/California/Arizona
Member Is Offline
Mood: decoherenced
|
|
Ah I see, I'll be sure to do more research on the topic. Thanks guys! (also i see now, earlier i was reading up on a dessicator bag and wondering how
sodium hydroxide 'absorbed' the moisture, i see now its a hydration type thing). Thanks again <3
-Capn' Rum the Science Bum
|
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
