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Home , Coordination complex

Coordination (see Chapter 20, H&S 3rd Ed.)

Transition elements, particularly , have a strong tendency to form compounds with Lewis bases. Why?

Lewis : electron pair donor egs. H2O, NH3 n+ Lewis acid: electron pair acceptor egs. M , BH3

‘low lying’ empty orbitals: lower in energy than HOMO of Lewis case

Lewis acidity enhanced by positive charge so TM ions are good Lewis bases

2+ 3+ familiar examples: AlCl3, BF3 from main and M , M in TM series

Type of bonding in these compounds?

‘coordinative’ or ‘dative’ bonds to form Lewis base ‘adducts’

L→M

TM-Lewis base ‘adducts’ are often referred to as (i) TM ‘complexes’ OR (ii) coordination complexes

NB: still polar covalent bonds – this just defines where the pair of electrons came from Some more definitions:

: or that can act as a Lewis base to a metal - egs. Cl , PR3, ROR (R = generic term for an alkyl) donor : atoms that are in direct contact with the metal ion O in R2O→M : number of donor atoms per ligand; range from 1 to >6

unidentate bidentate tridentate H N NH O H2O 2 2 O O water en diglyme

N O O O N N O O N

porphin 15-C-5 tetradentate pentadentate chelate: multidentate forms rings; M tend to be very stable, OO especially 5 and 6 member chelate rings

acac

‘bite’ or ‘’: a measure of the span or size of a chelate ligand; depends on M-L distance

M OO

acac

coordination complex MLn

generic terms commonly used:

L neutral Lewis base (electron pair donor)

X anionic donor (eg. Cl-)

κ indicates how many and which atoms donate in cases where it may be ambiguous

μ indicates ligand is bridging two (if no superscript) or more atoms

η or hapto number: indicates number of carbons directly bonded to the metal in an organometallic compound although you will see it incorrectly used for non-C coordination quite a bit

Alfred Werner (early 1900’s) developed what we now know as coordination theory to explain some odd observations: addition of NH3 to aqueous CoCl3 gave not one, but FOUR different complexes with three different empirical formulae!

CoCl3•4NH3 (• = ‘adduct’) one green complex one violet complex

CoCl3•5NH3 purple complex

CoCl3•6NH3 yellow complex

Why does it even react if it is neutral CoCl3? (eg. NH3 + CH4 = no rx)

What do you think is going on here and how might you test it (in 1910!)?

Werner used gravimetric precipitation to see whether the Cl- were ‘free’: precipitate Cl- as an insoluble salt, weigh it and see how many Cl- are in ionic form.

+ CoCl3•4NH3 + excess Ag → 1 AgCl for both

+ CoCl3•5NH3 + excess Ag → 2 AgCl

+ CoCl3•6NH3 + excess Ag → 3 AgCl

-10 + Ksp of AgCl is 1.8 x 10 (AgNO3 as source of soluble Ag ) Werner’s conclusions:

for each complex some Cl- are ‘free’ as anions

- - total number of NH3 and Cl is 6 in each case with ‘extra’ Cl present as free counterions to balance the charge

Werner therefore proposed that Co had: a ‘primary ’: of the metal ion a ‘secondary valence’: number of attached ligands (we now call this the or CN)

+ - i.e. CoCl3•4NH3 = [Co(NH3)4Cl2] Cl

2+ - CoCl3•5NH3 = [Co(NH3)5Cl] 2Cl

3+ - CoCl3•6NH3 = [Co(NH3)6] 3Cl

What is the dn number of Co in these complexes?

dn = 6 in all cases since these are all Co3+ Okay, but what are the actual structures? Werner had an answer for that too:

He focused his attention on the observation of two distinct + - structures with the formula [Co(NH3)4Cl2] Cl (green and violet):

three isomers Cl Cl Cl Cl hexagonal planar Co Co Co Cl Cl

three isomers Cl Cl Cl Cl trigonal prismatic Cl Cl

two isomers

Cl Cl octahedral Cl

Co Co

Cl Werner therefore concluded that these were Oh complexes with one, two or three ‘free’ Cl- anions (Nobel Prize 1913) The modern interpretation of Werner’s findings:

‘Inner coordination sphere’

ligands bound directly to the metal ion

‘Outer coordination sphere’

loosely associated counterions to balance charge (especially water) – note these can often be removed while ‘inner sphere’ waters cannot