Download The Transition Elements and Their Coordination Compounds

The Transition Elements and Their Coordination Compounds
Cu: plumbing
Fe: in steel
Cr: in automobile parts
Au, Ag: jewelry
W: light bulb filament
Ti: bicycle
Pt: auto catalytic converters
Zr: nuclear-reactor liners
Nitrinol (Ni and Ti used in stents)
And many more
Transition elements make up the d orbitals (we will cover here)
Inner transition metals make up the d and f orbitals
The transition elements (d block) and inner transition elements (f block) in the
periodic table
1
Example
Writing Electron Configurations of Transition Metal Atoms and Ions
Write condensed electron configurations for the following: (a) Zr; (b) V3+ (c) Mo3+.
(Assume that elements in higher periods behave like those in Period 4.)
Note that the general configuration is [noble gas] ns2(n - 1)dx. Recall that in ions the ns
electrons are lost first.
(a) Zr is the second element in the 4d series: [Kr] 5s24d2.
(b) V is the third element in the 3d series: [Ar] 4s23d3. In forming V3+, three electrons
are lost (two 4s and one 3d), so V3+ is a d2 ion: [Ar] 3d2
(c) Mo lies below Cr in Group 6B(6), so we expect the same exception as for Cr. Thus,
Mo is [Kr] 5s14d5. In forming the ion, Mo loses the one 5s and two of the 4d electrons,
so Mo3+ is a d3 ion: [Kr] 4d3
2
Horizontal trends in key atomic properties of the Period 4 elements
• Atomic size decreases overall across a period. The d electrons fill inner orbitals,
so they shield outer electrons from the increasing nuclear charge very efficiently
and the outer 4s electrons are not pulled closer.
• Electronegativity usually increases across a period but the transition metals
exhibit a relatively small change in electronegativity. Metal in higher oxidation
state is more positive has stronger pull on electron is more electronegative
• IE1 increase relatively little because the inner 3d electrons shield efficiently and
the outer 4s electron experiences only a slightly higher effective nuclear charge.
Vertical trends in key properties within the transition elements
3
Trends Within a Group (relative to main-group elements)
Atomic Size
Increases as expected, from Period 4 to 5. No increase from Period 5 to 6
Lanthanides with buried “4f” sublevel orbitals appear between the 4d (period 5) and 5d
(period 6) series
An element in Period 6 is separated from the one above it in Period 5 by 32 electrons
(ten 4d, six 5p, two 6s, and fourteen 4f). The extra shrinking that results from the
increased nuclear charge due to the addition of the fourteen 4f electrons is called the:
“Lanthanide Contraction”
Order of Sublevel Orbital Filling
Electronegativity (EN) – Relative ability of an atom in a covalent bond to attract shared
electrons.
EN of main-group elements decreases down group
Greater size means less attraction by nucleus
Greater Reactivity
EN in transition elements is opposite the trend in main-group elements
EN increases from period 4 to period 5.
4
No change from period 5 to period 6, since the change in volume is small and Zeff
increases (f orbital electrons).
Transition metals exhibit more covalent bonding and attract electrons more strongly
than main-group metals
The EN values in the heavy metals exceed those of most metalloids, forming salt-like
compounds, such as CsAu and the Au- ion
Ionization Energy
Main-group elements increase in size down a group, decreasing the 1st ionization
energy, making it relatively easier to remove the outer electrons
The relatively small increase in size of transition metals, combined with the relatively
large increase in nuclear charge (Zeff), results in a general increase in the first ionization
energy down a group
Density
Atomic size (volume) is inversely related to density
Across a period densities increase
In transition metals the density down a group increases dramatically because atomic
volumes change little from Period 5 to Period 6 while nuclear mass increases
significantly
Period 6 series contains some of the densest elements known:
Tungsten, Rhenium, Osmium, Iridium, Platinum, Gold (Density 20 times greater than
water, 2 times more dense than lead)
Oxidation States
5
• Exhibit more than one
• Ionic bonding is more prevalent for the lower oxidation states and covalent
bonding is more prevalent for the higher oxidation states.
At room temperature TiCl2 is an ionic solid and TiCl4 is a molecular liquid
• In high oxidation states atoms have higher charge densities ⇒ polarize electron
clouds of non-metals ⇒ covalent bonding
• The oxides become less basic as the oxidation state increases
TiO is a weak base in water and TiO2 is amphoteric.
Color and Magnetism of Compounds
Colors of representative compounds of the Period 4 transition
metals.
sodium
chromate
titanium oxide
scandium
oxide
nickel(II) nitrate
hexahydrate
potassium
zinc sulfate
ferricyanide
heptahydrate
manganese(II)
chloride
vanadyl sulfate tetrahydrate
dihydrate
cobalt(II)
chloride
hexahydrate
copper(II)
sulfate
pentahydrate
• Most main group ionic compounds are colorless because the metal ion has a
filled outer level;
On the contrary,
• Electrons in particular filled d-sublevels can absorb visible wavelengths and
move to slightly higher energy d-orbitals. Therefore, many transition metal
compounds have striking colors. Exceptions occur when d orbitals are empty
or filled. Zn2+: [Ar] 3d10 and Sc3+ or Ti4+ [Ar] 3d0
Formation of Coordination Compounds
These are species consisting of a central metal cation (transition metal or main group
metal) that is bonded to molecules and or anions called ligands. In order to maintain
neutrality in the coordination compound, the complex ion is typically associated with
other ions, called counterions.
Therefore,
6
Metals ions are Lewis acids, because they accept electrons from Lewis bases. When
metal cations combine with Lewis bases, the resulting species is called a complex ion,
and the base is called a ligand.
1s0 1s 11s 11s 1
+
H
H
H
+
+
N:
H
H
H
F
H
ammonia
sp3 orbitals
ammonium ion
H
N:
Hydrogen ion has a vacant 1s orbital.
Nitrogen has a full sp 3 orbital.
A co-ordinate covalent bond is formed
when nitrogen donates its lone e - pair
to hydrogen's empty 1s orbital.
N
H
H
ammonia
H
N
H
H
H H H
+
+
B
F
F
N
H
H
F
sp2 orbitals
B
F
F
B
N
sp3 orbitals
2pz
Boron has a vacant 2p z orbital.
Nitrogen has a full sp 3 orbital.
A co-ordinate covalent bond is formed
when nitrogen donates its lone e - pair
to boron's empty 2p z orbital.
Ammonia is the Lewis base (electron donor). H+ and BF3 are Lewis acids (electron
acceptors).
Similarly, transition metals and their cations are Lewis acids. They have vacant orbitals
that can accept electron pairs from donor atoms (Lewis bases) forming coordination
compounds or complexes. Some examples are shown in the following table.
Lewis Acid
Cr+3
+
Lewis Base
6H2O
→
Complex
[Cr(OH2)6]+3
Dissociation Constants
--
Co+3
+
6NH3
→
[Co(NH3)6]+3
2.2 × 10-34
Ni+2
Fe
Ag+
+
+
+
4CN5CO
2NH3
→
→
→
[Ni(CN)4]-2
[Fe(CO)5]
[Ag(NH3)2]+
1.0 × 10-31
-6.3 × 10-8
Lewis bases (anions or molecules) bonded to the central Lewis acid are called ligands
(Latin = tie or bind). Charged coordination complexes also have counter ions
associated with them to satisfy electroneutrality. For example, in [Ag(NH3)2]+NO3-, a
nitrate anion is the counter ion. Counter ions are ionically bonded. When dissolved in
water, the NO3- ion separates from the complex, but the two NH3 ligands remain firmly
bound to the Ag+ ion by covalent bonds.
7
Components of a Coordination Compound
When solid complex dissolves in water, the complex ion and the counter ions separate,
but ligands remain bound to central atom
Coordination Numbers, Geometries, and Ligands
Coordination Number (CN) - the number of ligand atoms that are bonded directly to
the central metal ion. The coordination number is specific for a given metal ion in a
particular oxidation state and compound (6 is the most common, 2 and 4 are often used.)
Geometry - the geometry (shape) of a complex ion depends on the coordination number
and nature of the metal ion. – See table below
Donor atoms per ligand - molecules and/or anions with one or more donor atoms that
each donate a lone pair of electrons to the metal ion to form a covalent bond.
Importance of Coordination Complexes:
Many biologically important substances are 'd'-transition metal coordination compounds
that are made up of large organic molecules bound to the metal via coordinate covalent
bonds, e.g.,
‰ hemoglobin (blood protein) is a coordination complex involving Fe.
‰ vitamin B-12 is a cobalt complex (cyanocobalamin)
‰ phthalocyaine blue is a Cu complex dye used for blue jeans and ink
‰ complexing agents are used for water softening (Ca, Mg, Fe removal with EDTA)
8
‰
‰
an antidote for some metal poisoning (BAL) forms complexes with As, Hg, and Cr.
Fe-carbonyls are anti-knock gasoline additives
Structure of Coordination Compounds:
Structures are determined by the coordination number (CN) and VSEPR. Lone pairs of
e-'s in d-orbitals have minimal influence on geometry because they are not in the outer
shell.
Geometries of Various Coordination Numbers
CN
Geometry
2
Hybridization
(nonmetals)
Hybridization
(transition
metals)
sp (BeH2)
sp
[Ag(NH3)2]+
[Cu(CN)2]-
sp3 (CH4)
sp3
[Zn(CN)4]-2
[Cd(NH3)4]+2
[FeCl4][Co(Br)4]-2
sp3d2 (XeF4)
dsp2
[Pd(CO)4]+2
[Pt(NH3)2Cl2]
[Ni(CN)4]-2
sp3d (PCl5)
d3sp or dsp3
[CuCl5]-3
[Fe(CO)5]
[Mn(CO)4NO]
[Ni(CN)5]-3
linear
4
109.5º
tetrahedral
4
90º
90º
Examples
(transition metals)
square planar
a
5
90º
or sp3d
e
e
120º
e
a
6
triangular bipyramidal
e = equitorial position
a = axial position
90º
90º
octahedral
sp3d2 (SF6)
d2sp3 or sp3d2
[PtCl6]-2
[Co(NH3)6]+3
[Fe(OH2)6]+2
[Fe(CN)6]-4
[MoF6][CoF6]-3
[Ni(H2O)6]+2
[Mn(CN)6]-4
CN 7 [UF7]-3, [ZrF7]-3, CN 8 [Mo(CN)8]-4, [TaF8]-3 and CN 9 [UCl3 hydrate] are known
but rare.
9
Typical CN for some common ions
M+
CN
M2+
CN
M3+
CN
+
+
2+
2+
2+
2+
2+
3+
3+
3+
Cu , Au 2,4 Co ,Ni , Cu ,Zn Mn 4,6 Sc , Cr , Co 6
Ag+
2
Fe2+
6
Au3+
4
Types of Ligands and Their Names
Rules for writing formulas for the
coordination compounds
1. The cation is written before the anion.
2. The charge of the cation(s) is balanced by
the charge of the anion(s).
3. For the complex ion, neutral ligands are
written before anionic ligands, and the
10
formula for the whole ion is placed in brackets.
Relative Strengths of Ligands:
Lewis bases differ in their ability to donate electrons to the metal ion. The relative
strengths of ligands in coordination complexes has been experimentally determined …
I-<Br-< Cl- < F- <OH-<C2O4-2<H2O< SCN-<NH3 <en < NO2-<CN-<CO
WEAK
STRONG
We will come back on this series.
Rules for naming complexes
1.
2.
3.
4.
5.
6.
The cation is named before the anion.
Within the complex ion, the ligands are named, in alphabetical order, before the
metal ion.
Neutral ligands generally have the molecule name, but there are a few exceptions.
Anionic ligands drop the -ide and add -o after the root name.
A numerical prefix indicates the number of ligands of a particular type.
The oxidation state of the central metal ion is given by a Roman numeral (in
parentheses) only if the metal ion can have more than one state.
If the complex ion is an anion, we drop the ending of the metal name and add -ate.
formula
:NH3
H2O
name
ammonia
water
:C≡O:
:PH3
carbon monoxide
phosphine
ligand name
formula
-
ammine
Cl
-
aqua
F
-
carbonyl
:C≡N:
phosphine
OH
-
name
ligand name
chloride
chloro
fluoride
fluoro
cyanide
cyano
hydroxide
hydroxo
-
:N=O
nitric oxide
nitrosyl
:NO2
nitrite
nitro (NO2-)
NO3-
nitrate
nitrato
:NO2-
nitrite
nitrito (ONO-)
NH2-
amide
amido
*SO4-2
sulfate
sulfato
thiocyanate
thiocyanato
-2
*C2O4
-2
*CO3
*O
-2
oxalate
oxalato
SCN
-2
carbonate
carbonato
*S2O3
thiosulfate
thiosulfato
oxide
oxo
C5H5N:
pyridine
pyridine
Example
a) What is the systematic name of Na3[AlF6]?
The complex ion is [AlF6]3-.
Six (hexa-) F- ions (fluoro) as ligands - hexafluoro
Aluminum is the central metal atom - aluminate
Aluminum has only the +3 ion, so we do not need Roman numerals.
∴ sodium hexafluoroaluminate
11
b) What is the systematic name of [Co(en)2Cl2]NO3?
There are two ligands, chlorine and ethylenediamine - dichloro, [bis(ethylenediamine)]
The complex is the cation and we have to use Roman numerals for the cobalt oxidation
state since it has more than one - (III)
The anion, nitrate, is named last.
∴ dichlorobis(ethylenediamine)cobalt(III) nitrate
c) What is the formula of tetraamminebromochloroplatinum(IV) chloride?
4 NH3,
Br − , Cl − , Pt4+
Cl −
∴ [Pt(NH3)4BrCl]Cl2
d) What is the formula of hexaamminecobalt(III) tetrachloro-ferrate(III)?
6 NH3
Co3+
4Cl −
Fe3+
∴[Co(NH3)6][FeCl4]3
Chelates
chelate with one
two
ethylenediamine ligand
three
Notice how the ligand “grabs” the metal from two sides like a claw
12
Coordination Complexes
Porphine
heme
Isomers (same chemical formula but different properties)
Coordination isomers occur when the composition of the complex ion changes but not
that of the compound
[Pt(NH3)4Cl2](NO2)2
and
[Pt(NH3)4 (NO2)2]Cl2
Linkage isomers occur when the composition of the complex ions remains the same but
the attachment of the ligand donor atom changes
[Co(NH3)5(NO2]Cl2 , [Co(NH3)5(ONO]Cl2
13
Geometric isomers (also called cis-trans and sometimes diastereomers) occur when
atoms or groups of atoms are arranged differently in space relative to the central metal
ion.
Beyond cis and trans isomers: facial & meridian isomers and enantiomers
CoCl3 ⋅3NH3
Cl
H3N Co Cl
H3N
Cl
NH3
facial (fac)
Cl
H3N Co Cl
H3 N
NH3
Cl
meridian (mer)
facial: 3 NH3 and 3 Cl ligands are adjacent (on triangular face)
meridian: 3 NH3 ligands in one plane, 3 Cl ligands in a perpendicular plane
Optical isomers (also called enantiomers) occur when a molecule and its mirror image
cannot be superimposed. Optical isomers have distinct physical properties like other
types of isomers, with one exception – the direction in which they rotate the plane of
polarized light.
14
Enantiomers: non superimposable mirror images
A structure is termed chiral if it is not superimposable on its mirror image
!
Two chiral structures: non superimposable mirror images: Enantiomers!
Enantiomers
NH3
H3N Co Cl
H2O
Cl
H2O
NH3
Cl Co NH3
Cl
H2O
H2O
15
Chirality: absence of a plane of symmetry, possible enantiomers
NH3
Cl Co H2O
Cl
H2O
NH3
Achiral (plane of symmetry)
NH3
H3N Co Cl
H2O
Cl
H2O
NH3
Cl Co NH3
Cl
H2O
H2O
No plane of symmetry: Chiral
Which are enantiomers (non-superimposable mirror images) and which are identical
(superimposable mirror images)?
16
- Rotating structure I in the cis compound gives structure III, which is not the same as
structure II, its mirror image, Image I & Image III are optical isomers
- Rotating structure I in the trans compound gives structure III,which is the same as
structure II, its mirror image, The trans compound does not have any mirror images
17
Determining the Type of Stereoisomerism
PROBLEM: Draw all stereoisomers for each of the following and state the type
of isomerism:
:
PLAN:
(b) [Cr(en)3]3+ (en = H2NCH2CH2 NH2)
:
(a) [Pt(NH3)2Br2]
Determine the geometry around each metal ion and the nature of
the ligands. Place the ligands in as many different positions as
possible. Look for cis-trans and optical isomers.
SOLUTION: (a) Pt(II) forms a square planar complex and there are two pair
of monodentate ligands - NH3 and Br.
These are geometric isomers;
they are not optical isomers
since they are superimposable
on their mirror images.
(b) Ethylenediamine is a bidentate ligand. Cr3+ has a coordination number of 6 and an
octahedral geometry.
Since all of the ligands are identical, there will be no geometric isomerism possible
The mirror images are nonsuperimposable and are therefore,
optical isomers.
18
Application of VB Theory to Complex Ions
Valence Bond Theory of Transition Metal Coordination Compounds:
VB theory treats metal to ligand (donor group) bonds as coordinate covalent bonds,
formed when a filled orbital of a donor atom overlaps with an empty hybrid orbital on
the central metal atom.
The molecular geometry is predicted using VSEPR.
The theory proposes that the number of metal-ion hybrid orbitals occupied by donor
atom lone pairs determine the geometry of the complexes. Lone pairs of electrons are
ignored. Since lone pairs are in the inner shell, they are believed to have little effect on
molecular geometry.
2-Coordinate Compounds:
Linear sp hybridized transition metal complexes:
a) Consider [Ag(NH3)2]+
5p
5s
4d
Ag+ = 5s0 4d10
Ag° = 5s1 4d10
sp
sp
+
[Ag(NH3)2]
Ag
5p
+
hybridized
:NH3 :NH3
Cu+1 = 4s0 3d10
b) Consider [Cu(CN)2]- Cu° = 4s1 3d10
4p
4s
3d
sp
sp
+
[Cu(CN)2]
Cu
4p
-
hybridized
+
K
CN CN
:C
N:
-
Four Coordinate Compounds: (Tetrahedral Complexes)
a) Consider [ZnCl4]-2
Zn+2 = 4s0 3d10
Zn° = 4s2 3d10
4p
4s
3d
sp3
sp3
+2
[ZnCl4]
Zn
-2
hybridized
tetrahedral
-
-
Cl Cl Cl Cl
b) Consider [Cd(NH3)4]+2
4d
5s
Cd° = 5s2 4d10
5p
sp3
sp3
+2
Cd
Cd+2 = 5s0 4d10
[Cd(NH3)4]
hybridized
tetrahedral
+2
..
..
.. ..
NH3 NH3 NH3 NH3
19
Note that all 4-coordinate complexes of a s0 d10 ion are sp3 hybridized and tetrahedral,
e.g.,
+2
‰ 4-coordinate complexes of Group 2B M
ions: Zn+2, Cd+2, Hg+2
+1
‰ 4-coordinate complexes of Group 1B M
ions: Cu+1, Ag+1, Au+1
In four coordinate complexes of s0d8 ions (Ni+2, Pd+2, Pt+2), two molecular geometries
are encountered, tetrahedral (sp3) and square planar (dsp2).
Consider [NiCl4]-2
3d
Ni° = 4s2 3d8
4s
4p
sp3
hybridized
+2
Ni
Ni+2 = 4s0 3d8
tetrahedral
[NiCl4]
-2
sp3
3d
As in previous examples of tetrahedral, sp3 hybridized complexes, the ligand donates
electrons to the vacant sp3 hybrid orbitals. This is true when large, weak ligands are
present. However, with small, strong ligands, such as CN-, two unpaired electrons in
half-filled 3d orbitals are forced to pair up with each other, creating an empty 3d orbital.
This gives rise to dsp2 hybridization and the square planar geometry seen in the
following examples.
Four Coordinate Compounds: (Square Planar Complexes):
When very strong ligands are present, four-coordinated d8 metal ions usually form
square planar complexes (rather than tetrahedral).
Consider [Ni(CN)4]-2
3d
Ni° = 4s2 3d8
4s
4p
Ni+2
3d
4s
4p
dsp2
hybridized [Ni(CN) ]-2
4
square
planar
Ni+2 = 4s0 3d8
dsp2
3d
-
4p
-
-
CN CN- CN CN
In dsp2 hybridization, one d-orbital is empty and receives an electron pair from the
ligand. Donor atoms are situated on the x and y -axes of the square planar structure.
Note that in the term 'dsp2', 'd' precedes 'sp2' indicating that the d-orbital used in
hybridization comes from a lower (inner) shell than the 's' and 'p' orbitals. This is called
inner shell hybridization. In the previous unit on bonding, we saw terms such as
'sp3d'. In this case, where 'd' follows 'sp3', the d-orbital is from the highest energy level,
e.g., recall that PCl5 has sp3d hybridization (triangular bipyramidal geometry). This is
called outer shell hybridization.
20
Six Coordinate Complexes (Octahedral):
Six ligand donor atoms in an octahedral complex are located at the corners of an
octahedron. Valence bond postulates that the orbitals best shaped and best oriented to
receive electron pairs from these directions are the dx2-y2 and dz2 orbitals (directed on the
x, y, and z-axes).
Consider [Fe(CN)6]-3
Fe° = 4s2 3d6
3d
Fe
4p
4s
+3
3d
Fe
Fe+3 = 4s0 3d5
4s
d2sp3
hybridized
4p
+3
octahedral
d2sp3
3d
[Fe(CN)6]
-3
-
-
-
-
CN CN- CN CN- CN CN
CN- is a strong field donor, forcing inner 3d electrons into a lower spin state and
hybridizing inner 3d orbitals, i.e., d2sp3 (inner shell hybridization). H2O, a weaker field
donor, is not able to bring about inner shell hybridization and the resulting hybridization
is thus sp3d2. Both d2sp3 and sp3d2 complexes are octahedral.
Consider [Fe(H2O)6]+3
3d
Fe
Fe° = 4s2 3d6
sp 3d2
4d
4p
4s
Fe+3 = 4s0 3d5
hybridized
+3
octahedral
sp 3d2
3d
[Fe(H2O)6]
4d
+3
H2O H2O H2O H2O H2O H2O
Metal ions in which two d-orbitals are vacant will form inner shell hybridized d2sp3
octahedral complexes.
Metal ions in which two d-orbitals are not vacant will form either outer shell sp3d2
hybridized octahedral complexes (with weak ligands) or inner shell d2sp3 hybridized
octahedral complexes with strong ligands (If half-filled d-orbitals are present a low spin
state can be forced).
Study the following examples and note where inner and outer shell hybridization occur
and why.
d6
x
x
Co+3
3d
[Co(NH3)6]
x
x
x
x
4d
4p
4s
x
x
+3
= an electron pair donated from a ligand
x
x
x
x
x
x
strong ligand
d2sp 3
3d
[CoF6]
-3
x
x
4d
4p
4s
x
x
x
x
x
x
sp 3d2
x
x
x
x
weak ligand
21
d5 Mn+2
3d
[Mn(CN)6]
-4
x
x
4d
4p
4s
x
x
x
x
x
x
x
x
x
x
d2sp 3
d8
Ni+2
3d
[Ni(H2O)6]
x
x
4d
4p
4s
+2
x
x
x
x
x
x
x
x
x
x
sp 3d2
d3
Cr
+3
3d
[Cr(NH3)6]
+3
x
x
x
x
4d
4p
4s
x
x
x
x
x
x
x
x
d2sp 3
What is color?
White light: EM radiation consisting of all λ’s in the visible range.
Objects appear colored in white light b/c they absorb certain λ’s and transmit (reflect)
other λ’s.
The transmitted light enters the eye, hits the retina and the brain perceives a color.
• If an object absorbs all λ’s ⇒ black
• If an object reflects all λ’s ⇒ white
• If an object absorbs all λ’s except for green , the reflected (transmitted) green
enters brain and it is interpreted as green.
• If an object absorbs only red (complementary of green) the remaining mixture
of reflected (transmitted) λ’s is entered the brain and it is interpreted as green.
22
Valence Bond Theory in Coordination Compounds: An Overview
ƒ Ligands (Lewis base) form coordinate covalent bonds with metal center (Lewis
acid)
ƒ Relationship between hybridization, geometry, and magnetism
ƒ Inadequate explanation for colors of complex ions
e.g., [Cr(H2O)6]3+, [Cr(H2O)4Cl2]+
Crystal Field Theory
Basis: purely electrostatic
Spherical Field: d orbitals degenerate
Overview of d orbitals
23
Splitting of d-orbital energies by an octahedral field of ligands
∆ is the splitting energy
Factors affecting magnitude of ∆
1. Oxidation state of the metal ion
2. Nature of the metal ion
3. Number and geometry of the ligands
4. Nature of the ligands
The effect of the ligand on splitting energy
24
The color of [Ti(H2O)6]3+
Effects of the metal oxidation state and of ligand identity on color
[V(H2O)6]3+
[V(H2O)6]2+
[Cr(NH3)6]3+
[Cr(NH3)5Cl ]2+
• For a given ligand, the color depends on the oxidation state of the metal ion.
• For a given metal ion, the color depends on the ligand.
The spectrochemical series
25
Ranking Crystal Field Splitting Energies for
Complex Ions of a Given Metal
PROBLEM: Rank the ions [Ti(H2O)6]3+, [Ti(NH3)6]3+, and [Ti(CN)6]3- in terms of
the relative value of ∆ and of the energy of visible light absorbed.
PLAN: The oxidation state of Ti is +3 in all of the complexes so we are
looking at the crystal field strength of the ligands. The stronger the
ligand, the greater the splitting, and the higher the energy of the
light absorbed.
SOLUTION:
The ligand field strength is CN- > NH3 > H2O, so the relative
size of ∆ and energy of light absorbed is
[Ti(CN)6]3- > [Ti(NH3)6]3+ > [Ti(H2O)6]3+
High-spin and low-spin complex ions of Mn2+
26
Identifying Complex Ions as High Spin or Low Spin
PROBLEM:
PLAN:
Iron (II) forms an essential complex in hemoglobin. For each of the
2+
4two octahedral complex ions [Fe(H2O)6] and [Fe(CN)6] , draw an
orbital splitting diagram, predict the number of unpaired electrons,
and identify the ion as low or high spin.
2+
The electron configuration of Fe gives us information that the iron
has 6d electrons. The two ligands have different field strengths.
Draw the orbital box diagrams, splitting the d orbitals into eg and
t2g. Add the electrons noting that a weak-field ligand gives the
maximum number of unpaired electrons and a high-spin complex
and vice-versa.
SOLUTION:
PE
[Fe(CN)6]4-
[Fe(H2O)6]2+
4 unpaired e
(high spin)
eg
t2g
--
eg
no unpaired e
(low spin)
--
t2g
Orbital occupancy for high-spin
and low-spin complexes of d4
through d7 metal ions
Splitting of d-orbital energies by a tetrahedral
field of ligands
Splitting of d-orbital energies by a square planar
field of ligands
27
Organometallic compounds: σ and π bonded compounds
Ni(CO)4 and Fe(CO)5
Bonding in carbonyls
Donation of lone pair of electrons on the carbonyl
carbon into a vacant orbital of the metal.
–
M
+
C
O:
σ-overlap
–
M
C
O:
donation of a pair of electrons from a filled d-orbital of
metal into vacant antibonding (π*) orbital of CO.
M
+
C
M
O:
C
O
π-overlap
28
Bonding in alkenes
π-electron density of the alkene overlaps with a σ-type
vacant orbital on the metal atom
C
+
M
C
σ - overlap
M
C
C
flow of electron density from a filled d-orbital on the metal into
the vacant p* antibonding molecular orbital on the carbon atoms
C
M
+
C
Back-bonding
C
M
C
29