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Transition metal bonding

Ionic bonding is a metal and a non metal, and covalent is 2 non metals.

How come transition metals form a co-ordinate bond between the ligand and the metal?

As the bonding should be ionic between a non-metal (ligand) and a metal (metal) and a co-ordinate bond is covalent.

(edited 13 years ago)

Reply 1

shengoc

Original post by uxa595

bonding is never really pure ionic or pure covalent in most compounds. The "rules" you stated were merely empirical and in simple main group compounds would apply.

Reply 2

uxa595

I don't get it though, how can a ligand and a metal form a covalent (co-ordinate) bond ?

The rules don't apply for transition metals/complex ions?

Reply 3

Tobia_s

The complex ions form because ligands are usually anions and have 1 or 2 lone pairs hanging about while transition metal ions are always cations. They are already slightly attracted to each other due to their being opposite in charge and the ligands then donate their lone pair(s) to the transition metal ion to form a complex ion.

(edited 13 years ago)

Reply 4

Tobia_s

Also; most ligands can also form ionic bonds with transition metal ions.

It could be a bit tricky to decide whether a 'species' is ionically bonded to the metal ion or whether it is acting as a ligand though.

Reply 5

uxa595

in the work i'm doing, ligands are bonded to metals using a co-ordinate bond.
I co-ordinate bond is always covalent.

So, i should just assume the rule for ionic and covalent doesn't apply?

Reply 6

illusionz

Firstly, don't think about ionic/covalent bonding in terms of metal/non-metal. Think about it in terms of electronegativity. A large electronegativity difference will mean an ionic bond, and a small difference will mean a covalent bond.

You need to appreciate that bonding is never purely ionic or purely covalent, there's a spectrum where you have NaCl at one end and methane at the other, but there are many compounds where it is not as clear cut.

Transition metals have partially filled d orbitals, and simply put, occupying them with electrons (from a ligand) can lower the overall energy of the species present, making it a favourable reaction. If you've learnt about molecular orbital/crystal field theory then you should know about this, if not, don't worry.

(edited 13 years ago)

Reply 7

shengoc

Original post by uxa595

in the work i'm doing, ligands are bonded to metals using a co-ordinate bond.
I co-ordinate bond is always covalent.

So, i should just assume the rule for ionic and covalent doesn't apply?

you should know by now, a lot of a level stuff are taken for granted. in chemistry alone, you can find lots of general rules, but it all comes down to some cases where minor tweaks are required.

bonding is a major topic in inorganic chemistry. bonding in compounds can stretch from pure ionic like LiF to really covalent like character, ie CsI; polarisation of electron clouds give rise to distorted hard spheres modelled by classic ionic bonding theory, hence giving rise to partial covalency.

Like the other poster said, bonding in complexes is with ligands. These ligands can be mono-(NH3), bi-(H2N-NH2) or polydentate(EDTA). They can be charged(ie halides) or not(NH3, water)

Again there are some really good books on bonding. but to simplify matters, these ligands can be categorised into sigma donor, pi-donor or pi-acceptor => these are covalent type of bonding(dative covalent/ hybridised covalent) and allows construction of molecular orbital diagrams to explain bonding; as for whether this is true or not, that is where photoelectron spectroscopy comes in and confirm the relative energies of these molecular orbitals involved in bonding.

But, we are not discounting ionic bonding as being important too. They play important role in solvation. Remember that most of these complex solutions are made in water. So this leads to another big chemistry field, thermodynamics in coordination chemistry.

I'll stop there. :biggrin: