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Displacement reaction

jamit - 21-6-2011 at 11:48

I was wondering why the displacement reaction happens. I know that activemetals replaces the lesser active metals. For example aluminum is more active than copper and so putting aluminum foil in copper II chloride will replace the copper and form aluminum chloride and copper.

Now the question I have is, will aluminum replace any copper bases salts, such as copper sulfate, copper nitrate, and any soluble copper salt? And if not why not? Can some one explain that to me or direct me in the right path? Thanks.

Mixell - 21-6-2011 at 11:57

Aluminium replaces the copper(II) ion:
2Al(s) +3Cu2+ --> 2Al3+ +3Cu(s).
So it will work with any soluble salt.

LanthanumK - 21-6-2011 at 12:16

Yes, the reaction works with all copper salts, but it is especially rapid in copper(II) chloride for reasons not exactly known.

blogfast25 - 21-6-2011 at 13:31

The term ‘active’ or ‘less active’ is very vague and non-scientific. Whether a metal will be able to displace another metal is quite dependent on conditions (solution, dry mix, state etc). In general for a chemical reaction be able to proceed, the Gibbs Free Energy of the end-state (G2

has to be lower than the starting state (G1

, so that ΔG = G2 - G1 < 0.

This is the case for the aqueous displacements you’re referring to.

What’s more, the copper salts used dissociate completely in water:

CuSO4 (s) (in water) === > Cu2+ (aq) + SO4(2-) (aq)

The sulphate (or chloride or nitrate or whatever) anion plays no part in the reaction:

3 Cu2+(aq) + 2 Al(s) === > 3 Cu(s) +2 Al3+(aq)

The anions are ‘spectator ions’: they don’t actually do anything other than fulfil the charge neutrality requirement.

LanthanumK - 21-6-2011 at 13:43

Complexes are formed with certain anions such as chloride, though, which affect the reaction.

jamit - 21-6-2011 at 23:04

Thank you all for your reply... very informative. Although LanthanumK's answer is quite obtuse... perhaps because i'm so ignorant of the meaning of Gibbs Free Energy of the end-state.

Hey Lanthanumk, does Gibb's Free Energy explain why the displacement reaction is so much more active in copper II chloride than in copper II sulfate or copper II nitrate?

Again thanks for some stimulating thought. now back to researching the meaning of Gibbs free energy.

LanthanumK - 22-6-2011 at 03:00

See this (http://woelen.homescience.net/science/chem/exps/cu+al/index....) for more details on why the reaction of Al with CuCl2 is more vigorous than the reaction with CuSO4.

blogfast25 - 22-6-2011 at 05:30

While Lant is correct about the formation of cupric chlorocuprate (CuCl42-

he draws the wrong conclusion about the effect it has on the Gibbs Free Energy change of the displacement reaction.

For 3 Cu2+(aq) + 2 Al(s) === > 3 Cu(s) + 2 Al3+(aq) (in the absence of chloride or any other copper complexing agent) the Gibbs Free Energy change ΔG (1) can be easily calculated with Nernst from the cell potential of the redox reaction.

But for the case

3 CuCl42-(aq) + 2 Al(s) === > 3 Cu(s) + 2 Al3+(aq) + 12 Cl-(aq)

… the ΔG (2) of CuCl42-(aq) === > Cu2+ (aq) + 4 Cl-(aq) has to be added to the ΔG (1). Because ΔG (2) is positive, the overall new ΔG becomes less negative. This applies also where other complexing agents (like NH3) ‘mask’ the naked Cu2+ ion: the complexes somewhat dampen the displacement reaction…

Jamit, if you don't get this you need to brush up on thermodynamics applied to chemical reactions, see in particular Laws of Nernst and Hess.

[Edited on 22-6-2011 by blogfast25]

PHILOU Zrealone - 22-6-2011 at 06:04

Beware that the resulting Aluminium salt may hydrolize immediately and free some acid...that may attack a little the ultrafine copper powder if the anion is an oxydizing one

So
Al(ONO2)3 + 3H2O <-=> Al(OH)3 + 3 HNO3
Cu + HNO3 --> NxOy + Cu(NO3)2 + H2O

Also the "self sustaining" reaction is not written:
AlCl3 + 3H2O <-=> Al(OH)3 + 3 HCl
3 HCl + Al --> 3/2H2 + AlCl3

This means that CuCl2/Al has three reaction motors...
1°) The electrochemical redoxpotential between Al/Cu couples.
2°) The spontaneous hydrolysis of AlCl3 generating HCl that chew through the oxyd layer and the metalic Al.
3°) The precipitation of unsoluble gelous Al(OH)3.

blogfast25 - 22-6-2011 at 06:22

Quote: Originally posted by PHILOU Zrealone

Beware that the resulting Aluminium salt may hydrolize immediately and free some acid...that may attack a little the ultrafine copper powder if the anion is an oxydizing one

So
Al(ONO2)3 + 3H2O <-=> Al(OH)3 + 3 HNO3
Cu + HNO3 --> NxOy + Cu(NO3)2 + H2O

Also the "self sustaining" reaction is not written:
AlCl3 + 3H2O <-=> Al(OH)3 + 3 HCl
3 HCl + Al --> 3/2H2 + AlCl3

This means that CuCl2/Al has three reaction motors...
1°) The electrochemical redoxpotential between Al/Cu couples.
2°) The spontaneous hydrolysis of AlCl3 generating HCl that chew through the oxyd layer and the metalic Al.
3°) The precipitation of unsoluble gelous Al(OH)3.

Oh dear. Neither harmed nor hindered by any real knowledge, PHILOU’s on the case (run to the hills!)…

Aluminium salts aren’t THAT prone to hydrolysis that they immediately fall apart to Al(OH)3, although most are slightly acidic when dissolved in pure water, due to:

[Al(H2O)6]3+ + H2O < === > [Al(OH)(H2O)5]2+ + H3O+

In theory [Al(H2O)6]3+ is a multi-protic acid, in reality dissociation steps beyond the one depicted are very weak. Quite strong solutions of typical Al salts can be prepared without even any cloudiness (from Al(OH)3). Very strong hot alum (K2SO4.Al2(SO4)3) can be prepared without appreciable hydrolysis of the Al3+ hydrate ion.

PHILOU Zrealone - 22-6-2011 at 06:41

I know not all Al salts do hydrolyze but some do and so I used the term "MAY" in my sentence...

"Beware that the resulting Aluminium salt MAY hydrolize immediately and free some acid."

Mixell - 22-6-2011 at 08:01

Aluminium salts tend to hydrolyze in serious amounts when they are anhydrous, for an example the AlCl3, when its anhydrous it reacts violently with water to produce aluminium hydroxide and HCl, when its hydrated, it behave just like blogfast mentioned.
In aqueous solution, any aluminium that displaced copper and goes into solution, immediately becomes the Al(H2O)3+ species by acquiring water molecules as ligands.

blogfast25 - 22-6-2011 at 12:00

Quote: Originally posted by Mixell

Aluminium salts tend to hydrolyze in serious amounts when they are anhydrous, for an example the AlCl3, when its anhydrous it reacts violently with water to produce aluminium hydroxide and HCl, [...]

Interesting point, Mixell but you could be generalising a little too much: AlCl3 is a bit of an odd one out because it’s on the borderline between ionic and covalent (see low MP/BP, dimerised vapour (Al2Cl6), strong Lewis acid). I wouldn’t vouch for anhydrous aluminium sulphate to behave that way but it would be hard find info confirming or infirming that. Alum can be ‘burnt’ to the ahydrous mixture of Al2(SO4)3 and K2SO4, so that could be a model to check the sulphate’s behaviour with.

Mixell - 22-6-2011 at 12:21

I meant those that hydrolyze, do hat when they are anhydrous.
I know that aluminium sulfate (and tried it myself) even when anhydrous will not hydrolyze vigorously.