Transcript Chemistry of Coordination Compounds

Chemistry of
Coordination Compounds
Brown, LeMay Ch 24
AP Chemistry
24.1: Structure of Complexes
Complex: species in which a central metal ion
(usually a transition metal) is bonded to a group of
surrounding molecules or ions
Coordination compound: compound that contains
a complex ion or ions.
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A coordination compound, or complex, consists of:
 Metal ion
 Acts as a Lewis acid (e- pair acceptor)
 Electrophile: species that is “e- poor” and seeks e(gets attacked by nucleophile)

Ligand or complexing agent: molecule or ion with a
lone pair of e- that bonds to a metal ion
 Acts as a Lewis base (e- pair donor)
 Coordinate covalent bond: special type of
covalent bond, between a metal ion and ligand
 Nucleophile: species that is “e- rich” and seeks an
e- poor area of a molecule (seeks an electrophile)
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Lewis Structures of common ligands
NH3
CN-
S2O32-
SCN-
H2O (not always included in formula, however)
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Complexation reactions



Ligand usually added “in excess” on AP exam
Usually result in color changes (colors
generally originate from e- transitions in a
partially filled d shell)
Change properties of metal ion
 Thermodynamic (DH, DS, DG)
 Electrochemical (Eº)
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
The golden-orange compound is CoCl3*6NH3
while the purple compound only has 5
ammonia molecules in the coordinated
compound. As shown in the ball-and-stick
model, the chlorides serve as counter ions to
the cobalt/ammonia coordiation complex in
the orange compound, while one of the
ammonia molecules is replaced by Cl in the
purple compound. In both cases, the
coordination geometry is octahedral around
Co.
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Notation
Write complexes in brackets, with charge on outside
Ex: Cu2+ (aq) + 4 NH3 (aq) → [Cu(NH3)4]2+ (aq)
H
|
Cu2+ (aq) + 4 :N ─ H (aq) →
|
H
:NH3

2+
Cu
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Coordination number


Number of positions where a ligand can bond.
 Similar to oxidation state
Each metal ion has a characteristic (i.e., typical)
coordination number, which can be predicted
according to crystal field theory.
 Ag+: coordination number = 2 (2 ligand bonding
positions); results in a linear complex
Ex: Ag+ (aq) + 2 NH3 (aq) → [Ag(NH3)2]+ (aq)
+
H
H
H
|
|
|
Ag+ (aq) + 2 :N ─ H (aq) → H─ N:Ag:N─H
|
|
|
H
H H
(aq)
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
Zn2+ & Cu2+: coordination number = 4; tetrahedral
complex
Ex: Zn2+ (aq) + 4 H2O (l) → [Zn(H2O)4]2+ (aq)

Pt2+: coordination number = 4; square planar
complex (d8 e- structure)
Ex: Pt2+ (aq) + 4 CN- (aq) → [Pt(CN)4]2- (aq)
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
Al3+, Cr3+, and Fe3+: coordination number = 6;
octahedral complex
Ex: Cr3+ (aq) + 5 NH3 (aq) + Cl- (aq) →
[Cr(NH3)5Cl]2+ (aq)
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Is dependent on:
 Charge of ligand:
Ni2+: 6 NH3 or 4 CN- (since CN- transfers more
negative charge)

Size of ligand:
Fe3+: 6 F- or 4 Cl- (larger ions take up more space)
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24.2: Chelates & Polydentate ligands

Ligands with more than one bonding position

Ethylenediamine (“en”, C2H4N2), or oxalate, C2O42-
Ex: Cr3+ (aq) + 3 C2O42- (aq) → [Cr(C2O4)3]3- (aq)
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24.3: Nomenclature
1. Name cation before anion; one or both may be a
complex. (Follow standard nomenclature for noncomplexes.)
2. Within each complex (neutral or ion), name all
ligands before the metal.
 Name ligands in alphabetical order; ignore
numerical prefixes when alphabetizing

If more than one of the same ligands is present,
use a numerical prefix: di, tri, tetra, penta, hexa,
…
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

Neutral ligands: use the name of the molecule
(with some exceptions)
NH3 ammineH2O aquaAnionic ligand: use suffix –o
Br- bromoCN- cyanoCl- chloroOH- hydroxo-
3. If the complex is an anion, use –ate suffix
 Record the oxidation number of the metal in
parentheses (if appropriate).
Ex: [Co(NH3)5Cl]Cl2 pentamminechlorocobalt (III) chloride
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Nomenclature practice
1. K4[Fe(CN)6] potassium hexacyanoferrate (II)
2. [Cr(NH3)4(H2O)CN]Cl2
tetrammineaquacyanochromium
(III) chloride
3. Na[Al(OH)4] sodium tetrahydroxoaluminate
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* 24.5: Color & Magnetism

Atoms or ions with a partially filled d-shell
usually exhibit color because the e- transitions
fall within the visible part of the EM spectrum.


Ex: transition metals such as Cu2+ (blue) and Fe3+
(orange)
Therefore, those with empty or filled d-shells
are usually colorless.

Ex: alkali & alkaline earth halides, Al3+
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* 24.6: Crystal Field Theory

Created to explain why transition metal ions in complexes
(having unfilled d-shells) are not necessarily paramagnetic.

With coordination bonding, valence d-orbitals are not truly
degenerate. Instead, they “split”.
 Some are lower in energy (more stable) and some higher.
http://scienceworld.wolfram.com/chemistry/CrystalFieldTheory.html
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

The gap between the higher and lower energy levels is called
the crystal-field splitting energy, which varies with each
ligand, yielding different E, (different l, different colors).
e- in an “unfilled” d-shell can actually be all paired (i.e.,
diamagnetic).
Ex: Co3+ (has 6 d e-)
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