Transcript Powerpoint - mvhs

Chemistry of
Coordination Compounds
Brown, LeMay Ch 24
AP Chemistry
Monta Vista High School
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.
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: metal-ligand bond
 Nucleophile: species that is “e- rich” and seeks an e- poor
area of a molecule (seeks an electrophile)
Lewis Structures of common ligands
NH3
S2O32-
CN-
SCN-
H2O (not always included in formula,
however)
Complexation reactions
Ligand usually added “in excess” on AP
 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)

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.
Notation
Write complexes in square brackets, with charge on
outside
2+
2+
2+
Ex: Cu (aq)
H + 4 NH3 (aq) → [Cu(NH3)4] (aq)
Cu2+ (aq)
|
+ 4 :N ─ H (aq) →
|
H
Cu
:NH3

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
+
+
[Ag(NH
H 3)2] (aq)
H
H
|
|
|
Ag+ (aq) + 2 :N ─ H (aq) → H─ N:Ag:N─H
(aq)
|
|
|
H
H H

Zn2+ & Cu2+: coordination number = 4; tetrahedral
complex
Ex: [Zn(H2O)4]2+ (aq)

Pt2+: coordination number = 4; square planar complex
(d8 e- structure)
Ex: [Pt(CN)4]2- (aq)

Al3+, Cr3+, and Fe3+: coordination number = 6;
octahedral complex
Ex: [Cr(NH3)5Cl]2+ (aq)
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)
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-
http://chemlabs.uoregon.edu/GeneralResources/models/bidentate.html
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
If more than one of the same ligand is present, use a numerical
prefix: di, tri, tetra, penta, hexa, …
Ignore numerical prefixes when alphabetizing.
3.
◦
Neutral ligands: use the name of the molecule (with
some exceptions)
NH3 ammineH2O aqua-
◦
Anionic ligand: use suffix –o
BrbromoCN- cyanoClchloroOH- hydroxo-
If the complex is an anion, use –ate suffix
◦ Record the oxidation number of the metal in
parentheses (if appropriate).
pentamminechlorocobalt (III) chloride
Ex: [Co(NH3)5Cl]Cl2
Nomenclature practice
1. K4[Fe(CN)6] potassium hexacyanoferrate
2. [Cr(NH3)4(H2O)CN]Cl2
tetrammineaquacyanochromium (III)
chloride
3. Na[Al(OH)4]
sodium tetrahydroxoaluminate
* 24.5: Color & Magnetism

Atoms or ions with a partially filled dshell usually exhibit color because the etransitions 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 dshells are usually colorless.
◦ Ex: alkali & alkaline earth halides, Al3+
* 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


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-)