William L Masterton
Cecile N. Hurley
http://academic.cengage.com/chemistry/masterton
Chapter 15
Complex Ions
Edward J. Neth • University of Connecticut
Outline
1. Composition of complex ions
2. Geometry of complex ions
3. Electronic structure of complex ions
4. Formation constants of complex ions
Compounds of Transition Metals
• Components
• Transition metal
• Associated ions and molecules
• Counter-ion
Components of Complex Ions and Compounds of
Them
•
•
•
•
Composition of complex ions
Geometry of complex ions
Electronic structure of central metal atom or ion
Equilibrium constant for the formation of the complex
ion
Composition of Complex Ions
• Cu2+ (aq) + 4NH3 (aq) ⇌ Cu(NH3)42+ (aq)
• When ammonia is added to a solution of aqueous
copper(II) ion, the color changes from pale blue to
deep blue
• The species that accounts for the color is a
complex made up of copper(II) ion and four
associated ammonia molecules
The Nature of the Complex Ion, Cu(NH3)42+
• Each ammonia molecule has a lone pair on the
nitrogen
• A bond forms between the ammonia and the
copper(II) ion
• Both electrons come from ammonia
• Result is a coordinate covalent bond
Terminology
• The species is referred to as a complex ion
• Central metal ion
• Small molecules or ions surround it
• These are called ligands
• The number of atoms bonded to the metal is the
coordination number
• This may or may not be equal to the number of ligands,
as we will see
Complex Ion or Compound?
• The species, Cu(NH3)42+, is clearly a complex ion;
because of the charge, it cannot be a compound
• Complex ions can associate with other ions to form
compounds:
• [Cu(NH3)4]F2
• Two fluorides are needed to balance the positive
charge on the complex
• The square brackets indicate the complex ion, i.e.,
everything inside is part of the complex ion
Cu(NH3)42+
• The central metal ion is Cu2+
• The ligands are the NH3 molecules
• The coordination number is 4
Figure 15.1
The Nature of the Metal Ion
• Complex ions form from
• Transition metals
• Main-group metals: Al, Sn, Pb
• Complex ions exist in aqueous solution
• Consider Zn(NO3)2
• In water, the ion exists as Zn(H2O)42+
Lewis Acid-Base Principles
• Lewis bases donate lone pairs
• Lewis acids accept lone pairs
The Nature of the Complex Ion
• Consider that hydrated metal cations are acidic
• Brønsted
• Zn(H2O)42+ ⇌ Zn(H2O)3(OH)+ (aq) + H+ (aq)
• Lewis
• The metal can accept a pair of electrons and is
therefore a Lewis acid
• Consider that the ligand possesses at least one lone
pair
• This makes the ligand a Lewis base
• The complex ion can be described as the product of
a Lewis acid-base reaction
Lewis and Brønsted Acids and Bases
• Lewis bases are also Brønsted bases
• Can accept an electron pair (or pairs)
• Lewis acids need not be Brønsted acids
• The Lewis model broadens the definition of an
acid
Charges of Complexes
• Charge of complex = oxidation number of central
metal + charges of ligands
• Platinum(II)
Table 15.1 – Complexes of Pt2+
Example 15.1
Chelating Agents
• Ligands
• Any molecule or anion with at least one unshared
pair of electrons
• Some ligands have more than one pair of unshared
electrons
• These ligands can coordinate using multiple pairs of
electrons
• Ligands that can form more than one bond to a
central metal ion are called chelates
Two Common Chelating Ligands
• Oxalate, C2O42• Ethylenediamine, H2N-CH2CH2-NH2
• Both of these ligands coordinate in two places per
ligand species
• Bidentate chelating ligand
Figure 15.2
Geometry of Complex Ions
• Review of Chapter 7
• Recall that the geometry of a molecule can be
determined by the way in which the central atom
coordinates with terminal atoms
• For a complex ion, the geometry is determined by
the shape taken by the complex, as determined by
the coordination number and nature of the metal
ion and ligands
Coordination Number
• The most common coordination number is 6
• Geometry is octahedral
• Two other coordination numbers are common
• 2; Geometry is linear
• 4; Geometry may be square planar or
tetrahedral
Table 15.2
Figure 15.2
Example 15.2
Coordination Number 2
• 2-coordinate complexes are linear
• CuCl2• Ag(NH3)2+
• Au(CN)2-
Coordination Number 4
• Two geometries exist
• Tetrahedral
• Zn(NH3)42+
• CoCl42-
• Square planar
• Characteristic of Cu2+ and metal ions with d8
configurations (Pt2+, Ni2+)
Figure 15.3
Coordination Number 6
• 6-coordinate complexes are octahedral
• The six ligands are equidistant from the central
metal
• The octahedron can be considered a derivative of
a square plane, with two ligands added, one
above and one below the plane
Example 15.3
Isomerism
• Two or more species with the same formula, but
different chemical and physical properties are called
isomers
• Complex ions can show several different types of
isomerism
• Only type to be considered here is geometric
isomerism
• Only the spatial orientation of ligands differs
between geometric isomers
Isomerism in Square Planar Complexes
• Pt(NH3)2Cl2
• Can have two different isomers: cis and trans
• Complexes of the form Ma2b2 will show cis-trans
isomerism
• a and b are different ligands
Meaning of cis and trans
• Cis positions are 90° apart
• Trans positions are 180° apart
Figure 15.4
Isomerism in Octahedral Complexes
Co(NH3)4Cl2 or Ma4b2
Chelates and Isomers
• In general, chelating ligands can bridge only cis
positions
• The bridge is not long enough to stretch across a
trans position
• Chelates, due to their binding in two locations,
generally produce more stable complexes
• Partially a result of the geometry (ring size)
• Partially as a result of the nature of the bonds
Figure 15.6
Figure 15.7
Physical Properties of Isomers
• Note in the last figure that the color of the two
complexes differ
• The cis isomer is reddish-purple
• The trans isomer is green
• Geometric isomerism can lead to great differences in
the physical and chemical properties of compounds
containing complex ions
Electronic Structure of Complex Ions
• Crystal Field Model explains
• Color of transition metal complexes
• Magnetic properties of transition metal complexes
• Considers the nature of the
• Metal
• Ligands
Transition Metal Cations
• In a simple transition metal cation
• There are no outer s electrons
• Electrons are distributed among the five d orbitals
by Hund’s Rule
• Recall that Hund’s rule results in the maximum number
of unpaired spins
• Magnetic properties depend on distribution of electrons
• Diamagnetism: no unpaired electrons
• Paramagnetism: unpaired electrons
Iron(II) ion
• Iron(II) ion has 26-2 or 24 electrons
• Shorthand notation: [Ar]3d6
• Orbital diagram:
• [Ar] (↑↓) (↑ ) (↑ ) (↑ ) (↑ )
• Fe2+ is paramagnetic; there are four unpaired
electrons
Figure 15.8 – Colors of Transition Metal
Compounds
Example 15.4
d-orbitals
• Recall that there are five d orbitals
dz2 ,dx 2 y 2 ,dxy ,dxz ,dyz
• In uncomplexed metal cations, these orbitals have the same
energy
Figure 15.9
Octahedral Complexes
• We can collect the d orbitals into two groups:
• A high energy pair, the dx2-y2 and dz2
• A low energy triplet, the dxy, dxz and dyz
Why the Split into Two Groups?
• In the absence of any ligands, the d orbitals have the
same energy
• Note that the two orbitals whose energy is
considered high lie on the xyz axes
• Electron density in metal interacts with the
electron density of ligands brought toward the
metal on the axes
• Note that the two orbitals whose energy is
considered lower lie between the xyz axes
• There is less interaction between the metal
electron density and that of the ligand
The Splitting Energy, Δo
• The crystal field splitting energy is given the symbol
Δo
• The magnitude of Δo will determine the way in which
the electrons fill the d orbitals in the metal ion of the
complex
• If Δo is large, electrons will pair in the lower energy
orbitals before occupying the higher energy ones
• If Δo is small, electrons will distribute themselves
in all five orbitals, pairing only in cases where
there are more than five electrons
Figure 15.10
Figure 15.11
Electronic Arrangements in Complexes
• When Δo is large, a low spin complex results
• Electrons fill the lower three orbitals before
occupying the upper two
• When Δo is small, a high spin complex results
• Electrons distribute themselves to all five orbitals
by Hund’s rule
Notes on High and Low Spin
• For a given cation, a high spin complex will always
have a larger number of unpaired electrons than a
low spin complex
• The value of Δo is determined by the nature of the
ligand(s)
• Strong field-ligands produce low spin complexes
• Example: CN-, NH3
• Weak field-ligands produce high spin complexes
• Example: H2O, Cl-
Example 15.5
Color
• Most transition metal complexes are brightly colored
• Exception: those with empty d sublevels (e.g.,
Sc3+; those with full d sublevels (e.g., Zn2+)
• The splitting of the d sublevel results in an energy
difference that corresponds to the visible region of
the electromagnetic spectrum
• Visible light is absorbed in the transition of an electron
in the ion
• Some of the wavelengths of white light are removed
• Complex appears colored
Figure 15.12
Titanium(III)
• Consider Ti3+ in a complex
• There is only one d electron
• Because there is only one possible electronic
transition, it is possible to calculate Δo for this ion
• The ion absorbs at 510 nm (green)
• The complex appears as the color complement of green
(i.e., red-violet or purple)
Calculating Δo for Titanium(III)
E
hc

 3.90  10 19 J
23
1
kJ
6
.
022
X
10
E  3.90  10 19 J 

1000J
mol
• By knowing the energy of the light absorbed, it is
possible to calculate the value of Δo
• For ions with more than one d electron, calculating
the energy Δo is more difficult
Δo and Wavelength
• The smaller Δo is, the longer the wavelength of
absorption
• Small Δo stem from weak field ligands
• Large Δo stem from strong field ligands
Light Absorption and Color
The Spectrochemical Series of Ligands
• CN- > NO2- > en > NH3 > NCS- > H2O > F- > ClStrong field
weak field
Equilibrium and Complex Ions
• We can consider the formation of a complex ion in
aqueous solution as an equilibrium between
• The reactants
• Bare metal ion and ligands
• The product
• The complex ion
• As with acid-base and gas phase equilibria, an
expression for K can be written
Formation Constants
• Consider
• Cu2+ (aq) + 4NH3 (aq) ⇌ Cu(NH3)42+ (aq)
• We can write the expression for Kf (the formation constant for
the ion:
2
[Cu(NH3 )4 ]
Kf 
[Cu2 ][NH3 ]4
• In many cases, Kf is a very large number, indicating that the
complex which forms is very stable
Table 15.4
Example 15.6
Example 15.6 (Cont’d)
Practical Example
• Most cleaning compounds containing ammonia state
that the product is not to be used on objects made of
silver
• Consider the Kf for Ag(NH3)2+: 1.7 X 107
• The complex of silver with ammonia is a very
stable one
• Using an ammonia-based cleaner on silver will
partially dissolve the silver
Chelates in Nature
• Chelating agents are abundant in nature
• Iron is complexed with the chelate heme
• EDTA is used to complex metal ions
• Treat heavy metal poisoning
• Mask metal ions that promote food spoilage
Heme
EDTA
Key Concepts
1. Relate the composition of a complex ion to its
charge, coordination number, and the oxidation
number of the central metal
2. Sketch the geometry of a complex ion and identify
geometric isomers
3. Give the electron configuration and the orbital
diagram for the transition metal at the center of the
complex
4. Derive the orbital diagram (high or low spin) for the
complex
5. Relate the formation equilibrium constant to the
concentrations of metal ion, ligands and complex ion