Transcript Chapter 26 Transitional Metal complex (notes)

Chemistry 223 Chapter 26
Coordination Complexes
d-block elements a.k.a. transition metals
d-block elements are:
• all metals
• all have partially filled d subshells
• exhibit horizontal & vertical similarities
• alloys & compounds are important components
of materials in modern world
• most first-row transition metals are essential
for life
General Trends among Transition Metals
General Trends among Transition Metals
4th row Horizontal Periodic Trends
General Trends among Transition Metals
Going across row from left to right ,
e-’s are added to 3d subshell
to neutralize increase in
(+) charge of nucleus
as atomic # increases.
General Trends among Transition Metals
3d subshell fill based on
aufbau principle & Hund’s rule
with two important exceptions:
Reactivity:
Size of neutral atoms of d-block elements
gradually decreases
left to right across a row.
Why?
Due to increase in Zeff with increasing atomic #
Atomic radius increases going down a column.
Why?
Transition metals become less reactive (more “Noble”)
going from left to right across a row
Trends in Transition Metal Oxidation States:
Transition metals form cations by initial loss of ns e-’s,
even though ns orbital is lower in energy
than (n–1)d subshell in the neutral atom.
d-electron configuration for di-cations of
1st row of transition metals
Trends in Transition Metal Oxidation States:
Small E difference btwn ns and (n-1)d
plus screening effect means
less E losing ns e-’s before (n-1)d e-’s
All transition-metal cations possess dn
valence e- configurations for 2+ ions of 1st row.
Trends in Transition Metal Oxidation States:
Electronegativities of first-row transition metals
increase (somewhat) smoothly from Sc to Cu
Sc
1.36
Ti
1.54
V
1.63
Cr
1.66
Mn
1.55
Fe
1.83
Co
1.88
Ni
1.91
Cu
1.90
Zn
1.65
Trends in Transition Metal Oxidation States:
max oxid states for 2nd & 3rd row transition metals
in Groups 3 thru 8
increase from +3 for Y and La
to +8 for Ru and Os
Trends in Transition Metal Oxidation States:
Going farther to right,
maximum oxidation state decreases,
reaching +2 for elements of Group 12,
Descriptive Chemistry of 3d Transition Metals:
Scandium
Titanium
[Ar]4s23d1
[Ar]4s23d2
+3
+4
Vanadium
[Ar]4s23d3
+2, +3, +4, +5
Catalysts, steel alloys
Chromium
[Ar]4s13d5
+2, +3, +6
Colorful, Cr2O72− OA, stainless steel, chrome plating
Manganese
[Ar]4s23d5
+2, +4, +7
MnO4− OA, MnO2 catalyst, Mn steels
Iron
[Ar]4s23d6
+2, +3
Ores are hematite, magnetite, and pyrite (fool’s
gold), steel, hemoglobin, blast furnace, magnetic
Cobalt
Nickel
Copper
[Ar]4s23d7
[Ar]4s23d8
[Ar]4s13d10
+2, +3
+2
+1, +2
Blue cobalt glass, , AlNiCo, magnetic
Coins, AlNiCo, Monel, magnetic
Coins, brass, bronze, Statue of Liberty, patina,
electric wires, ores are chalcocite, chalcopyrite and
malachite, unreactive w/ HCl and H2SO4 but very
reactive w/HNO3
Zinc
[Ar]4s23d10
+2
Coins, brass, biochemistry, RA
Gold
[Xe]6s14f145d10
+1, +3
Coins, jewelry, soft as pure metal, alloys are
harder, CN− used to extract Au from ores
Silver
[Kr]5s14d10
+1
Coins, jewelry, most electrically conductive of all
metals
Mercury
[Xe]6s24f145d10
+1, +2
Quicksilver, poisonous, “mad as a hatter”,
Minimata
Strong, light, corrosion-resistant, steel alloys, white
pigments, ore is rutile
Compounds of Mn in +2 to +7 oxidation states
Different # of d electrons = different colors
Why is that?
Coordination Compounds
Metallic elements act as Lewis acids
form complexes with various Lewis bases.
Metal complex:
Coordination Compounds
Central metal atom (or ion) bonded to one or
more ligands.
Ligands:
Ligands
Coordination Compounds
metal & ligand complexes as ions:
Coordination Compounds
Coordination compounds & complexes are
distinct chemical species
properties & behavior diff from
metal atom / ion or the ligands
History of Coordination Compounds
Coordination compounds used since
ancient times, but chemical nature unclear.
Werner: modern theory of coordination
chemistry - based on studies of several
series of metal halide complexes with
ammonia
History of Coordination Compounds
Werner postulated that metal ions have
2 different kinds of valence:
1. primary valence (oxidation state) =
2. secondary valence (coordination #)
Alfred Werner
(1866-1919)
Same chemical composition, same # of groups
of same types attached to same metal.
What made the two different colors?
More on this later!
Structures of Metal Complexes
Coordination #’s of metal ions
in metal complexes
can range from 2 to 9.
Differences in E btwn
different arrangements of ligands greatest
for complexes w/ low coordination #’s
& decrease as coordination # increases.
Structures of Metal Complexes
Only one or two structures possible for
complexes w/ low coordination #’s.
Several different energetically = structures
are possible for complexes with
high coordination #’s (n > 6)
Structures of Metal Complexes
Coordination # 2 = linear
Rare for most metals;
common for d10 metal ions,
especially: Cu+, Ag+, Au+, and Hg2+
Coordination # 4
Two common structures:
tetrahedral & square planar
Tetrahedral: all 4-coordinate complexes of
• non-transition metals &
• d10 ions and first-row transition metals,
Coordination # 4
Two common structures: tetrahedral & square planar
Square planar: 4-coordinate complexes of
2nd & 3rd row transition metals
with d8 e- configurations,
e.g. Rh+ , Pt2+ and Pd2+,
also encountered in some Ni2+ & Cu2+ complexes.
Structures of Metal Complexes
Coordination # 6
Most common: six ligands at vertices of an
octahedron or a distorted octahedron.
We will focus primarily on octahedral
Other Structures of Metal Complexes Possible:
Coordination # 3
Encountered with d10 metal ions e.g.Cu+ & Hg2+
trigonal planar structure
Coordination # 5 geometries
… and 2nd & 3rd row
transition metals
7, 8 & 9
coordination #’s,
give other geometries:
Metal-ligand interaction is an example of
Lewis acid-base interaction.
Lewis acid
Lewis base
Lewis bases
Must have
Transition metal ions tend to form
coordination complexes
which we encountered back in Chapter 22.
e.g. AgCl is more soluble in 0.10 M NH3
than it is in pure water because
Ag+ forms a complex with NH3
with a very large formation constant:
Ag+ + 2NH3  Ag(NH3)2+
The complex ion Ag(NH3)2+
that forms is called diamminesilver(I)
(review rules on pp. 1055-1056).
Why does it form?
It forms because each NH3 is a Lewis base
and forms a coordinate covalent bond
with the silver ion, Ag+, in solution
The complex has a linear geometry.
to purify the Ag(NH3)2+ complex ion
& store it in a bottle
it would need an anion
to neutralize the charge
e.g.
diamminesilver(I) chloride, [Ag(NH3)2+]Cl
or
diamminesilver(I) nitrate: [Ag(NH3)2+]NO3.
[Ag(NH3)2+]Cl or
[Ag(NH3)2+]NO3.
In these compounds,
silver is ____________
NH3 is ______________
and Cl or NO3 is ____________________.
Ligands are attached by ___________ bonds
Counterions are attached by _______ bonds!
Another complex formation reaction is:
Co3+ + 6 NH3  Co(NH3)63+
Kf = [Co(NH3)63+] = 2.3 x 1033
[Co3+][NH3]6
This complex ion is called:
This complex has
an octahedral geometry.
Another example is:
Cu2+ + 4 CN  Cu(CN)42
Kf = [Cu(CN)42] = 1.0 x 1025
[Cu2+][CN]4
This complex ion is called
This complex has
a tetrahedral geometry.
When a bidentate ligand binds to a metal,
A polydentate ligand is a chelating agent,
complexes containing polydentate ligands:
Ethylenediaminetetraacetate ion:
hexadentate ligand
chelate effect:
metal complexes
of
polydentate ligands
are more stable
than complexes
of
chemically similar
monodentate
ligands.
Nomenclature (IUPAC) rules for
Naming coordination compounds:
• Cation named before anion (as usual);
but, transition metal atom in the complex
is named last
with oxidation state in
roman numerals in parentheses
Nomenclature (IUPAC) rules for
Naming coordination compounds:
• Cation named before anion (as usual), no D;
• anion ending for transition metal will be “ate”
e.g. Cobalt anion =
[Ni(NH3)6] (NO3)2 cation complex
K3 [Co(Cl)6] anion complex
Anionic complex metal ending:
Scandium = Scandate
Titanium =
Titanate
Vanadium = Vanadate
Chromium = Chromate
Manganese = Manganate
Iron =
Ferrate
Cobalt =
Cobaltate
Nickel =
Nickelate
Copper =
Cuprate
Zinc =
Zincate
Special names for some transition metals
in an anion complex
Nomenclature (IUPAC) rules for
Naming complexes:
Ligands named 1st (alphabetically)
• Greek prefixes for counting
di, tri, tetra, penta, hexa, etc.
Nomenclature (IUPAC) rules for
Naming anionic ligands:
• Use suffix “o” if ending in “ide”
(e.g. chloride  chloro; cyanide  cyano
hydroxide  hydroxo; oxide  oxo)
• Use suffix “ito” if ending in “ite”
(e.g. nitrite  nitrito)
• Use suffix “ato” if ending in “ate”
(e.g. oxalate  oxalato; sulfate  sulfato
carbonate  carbonato
Neutral ligands:
Usual name: e.g. ethylenediamine
Exceptions:
Nitrite, NO2:
Which atom on the ligand
donates its lone pair D’s the name
Give the chemical formula for
Hexaaquanickel(ll) diaquatetrabromochromate(lll)
Give the chemical formula for
Give the name for
[Co(NH3)6][CoCl6]
Practice naming some complex compounds:
[Pt(Cl2)(NH3)2]
K2[PtCl4]
Practice naming some complex compounds:
[Pt(NH3)4]Cl2
[Pt(NH3)3Cl]Cl
Na[CoCl4(NH3)2]
Practice writing the
complex compound formulas:
hexaaquochromium(III) chloride
diaquodichloroaurate(III) chloride
potassium hexacyanoferrate(II)
potassium hexacyanoferrate (III)
Clicker Qstn:
the correct name for the complex
Na2[Ni(CN)4]
A. Disodium tetranickelcyanide
B. Sodium tetracyanidenickel(l)
C. Disodium tetracyanonickelo(lV)
D. Natrium tetranickel(Vl)cyanide
E. Sodium tetracyanonickelate(lll)
Constitutional (Structural) Isomers
1. Ionization isomers
2. Linkage isomers
Geometrical isomers of Complexes
Differ only in arrangement of ligands around metal ion.
Metal complexes that differ only in which ligands are
adjacent to one another (cis)
or
directly across from one another (trans).
Cis-platin isomer fights cancer,
Trans-platin doesn’t
Geometrical isomers are most important
for square planar & octahedral
complexes.
Square planar complexes:
all vertices of a square are equivalent,
it does not matter which vertex
is occupied by ligand B
in a square planar MA3B complex.
Only one geometrical isomer is possible
Only one isomer when there’s one B ligand.
With two, there are other possible arrangements.
Symmetrical bidentate ligands
also only have one structure
Isomers of Metal Complexes
Octahedral complexes:
Only one structure possible for octahedral complexes
(if only one ligand is different from other five):
(MA5B)
since all six vertices of an octahedron are equivalent.
Isomers of Metal Complexes
Octahedral complexes:
If two ligands in an octahedral complex
are different from other four (MA4B2),
two isomers are possible:
two B ligands can be _____________________.
Chelating agents (chelate = ________)
Bidentate (2 teeth):
Bidentate (2 teeth): e.g. ethylenediamine (en)
Octahedral isomer complexes:
Replacing another A ligand by B
gives an MA3B3 complex
for which there are two isomers:
Octahedral isomer complexes:
Fac: 3 ligands of each kind
occupy opposite triangular faces
of the octahedron
Mer: 3 ligands of each kind
lie on the meridian
(cut across flat mid-plane)
(cut across flat mid-plane)
Some coordination complexes with mixed ligands
have optical isomers and are said to be chiral.
A complex is chiral if its mirror images
are different molecules.
Anything that’s linear is not chiral (achiral),
i.e. mirror image is always same as original.
Anything that’s square planar is not chiral
(achiral),
i.e. mirror image is always same as original.
Anything that’s tetrahedral is chiral
only if all four groups are different.
octahedral is chiral with monodentate groups only if:
(a) all six groups are different (ABCDEF) or
(b) two groups are the same and cis (AABCDE) or
(c) three groups are the same and fac,
i.e. none trans (AAABCD) or
(d) there are two pairs of identical groups and both
are cis (AABBCD)
OR
some possibilities with bidentate ligands,
cis-dichlorobisethylenediaminecobalt(III)
If any pair of identical groups is trans,
there is no chirality!
any octahedral molecule with a mirror plane is achiral.
(any single pair of identical trans ligands
guarantees a mirror plane)
Crystal Field Theory
Crystal Field Theory
Bonding model explaining many important
properties of
transition-metal complexes:
Crystal Field Theory
Central assumption of CFT:
metal-ligand connections are
electrostatic interactions btwn
a central metal ion
and a set of negatively charged
ligands (or ligand dipoles)
arranged around metal ion.
d-Orbital Splittings
five d orbitals are initially degenerate
(same energy).
When the 6 (-) charges
are distributed uniformly over
surface of a sphere,
d orbitals remain degenerate.
d-Orbital Splittings
But!
Their energy will be higher
due to
d-Orbital Splittings
If the 6 (-) charges are placed at vertices
of an octahedron,
avg energy of d orbitals
does not change.
d-Orbital Splittings
But!
It does remove their degeneracy
and the 5 d orbitals split into two groups
d-Orbital Splittings
The dx2 – y2 and dz2 orbitals (eg orbitals)
point directly at the six (-) charges,
which increase their Energy compared with
a spherical distribution of negative charge.
The dxy, dxz, & dyz (t2g orbitals)
are all oriented at a 45º angle to the coordinate axes
and point between the 6 (-) charges,
which decreases their Energy
compared with a spherical distribution of charge
As LP’s onAs
ligands approach
along x, y, and z axes.
d-Orbital Splittings
Difference in E btwn the two sets of d orbitals
is
crystal field splitting energy.
d-Orbital Splittings
Magnitude of the splitting depends on:
Splitting of d orbitals in a crystal field
does not D total energy
of the five d orbitals
Electronic Structures of Metal Complexes
Using d-orbital energy-level diagram:
electronic structures & some properties of
transition-metal complexes can be predicted.
Start with Ti3+ ion,
(contains a single d electron),
proceed across first row of transition metals
by adding a single e- at a time.
Additional e-’s placed in lowest-E orbital available
while keeping their spins parallel
For d1-d3 systems,
e-’s successively occupy the 3 degenerate t2g orbitals
with their spins parallel (paramagnetic)
giving one, two, and three unpaired electrons.
Electronic Structures of Metal Complexes
d4 configuration: two possible choices for 4th e-:
enter one of the empty eg orbitals or
enter one of the singly occupied t2g orbitals
D<P
D>P
Spin Pairing Energy (P) is an increase in Energy
(due to electrostatic repulsions)
when an e- is put into an occupied orbital.
If Do is < P,
then lowest-energy arrangement has 4th ein an empty eg orbital.
Electronic Structures of Metal Complexes
If Do is > P,
lowest-energy arrangement has 4th ein one of the occupied t2g orbitals,
Metal ions with d4, d5, d6, or d7 e- configurations
can be either high spin or low spin,
depending on magnitude of Do
magnitude of Do
Large Do =
Smaller Do =
Only one arrangement of d electrons is possible
for metal ions with d8–d10 e- configurations
Factors That Affect the Magnitude of Do
magnitude of Do dictates whether
a complex with 4, 5, 6, or 7 d e-’s
is high spin or low spin:
1. Large values of Do yield a low-spin complex
2. Small values of Do  a high-spin complex
Which affects its:
• Magnetic properties
• Structure
• Reactivity
Factors That Affect the Magnitude of Do
Nature of the ligands
For a series of chem similar ligands,
magnitude of Do decreases
as size of donor atom increases
because smaller, more localized charges
interact
Factors That Affect the Magnitude of Do
Nature of the ligands
Nature of the ligands
experimentally observed order of
the crystal field splitting energies
produced by different ligands is called:
the spectrochemical series
Nature of the ligands
1. Strong-field ligands interact strongly
with d orbitals of metal ions
and give a large Do
2. Weak-field ligands interact
more weakly and give a smaller Do
The Spectrochemical Series
splitting of d orbitals in crystal field model
not only depends on geometry of the complex
also depends on nature of the metal ion,
charge on this ion,
and the ligands that surround the metal.
The Spectrochemical Series
When geometry and ligands are held constant,
this splitting decreases in the following order:
Pt4+ >
Ir3+ > Rh3+ > Co3+ >
strong-field ions
Cr3+ >
Fe3+ >
Fe2+ > Co2+ >
Ni2+ > Mn2+
weak-field ions
Metal ions at one end are called strong-field ions,
because splitting due to crystal field is
unusually strong.
Ions at other end are known as weak-field ions.
The Spectrochemical Series
When geometry & the metal are held constant,
splitting of d orbitals decreases in the following order:
CO
CN- > NO2- > NH3 > -NCS- > H2O >
strong-field
ligands
OH-
F-
-SCN-
Cl- >
Br-
weak-field ligands
Strong Field Ligands:
(strongest) CN−, CO > NO2− > en > NH3
Weak Field Ligands:
H2O > ox > OH− > F− > SCN−, Cl− > Br− > I− (weakest)
tetrahedral crystal field:
imagine 3 ligands lying at alternating corners of a cube
The dx2-y2 & dz2 orbitals
on metal ion at center of the cube
lie between the ligands,
and dxy, dxz, & dyz orbitals point toward the ligands.
Tetrahedral Complexes
Splitting of energies of orbitals in tetrahedral
complex, Do, is smaller than in an octahedral
complex for two reasons:
1. d orbitals interact less strongly with ligands
in a tetrahedral arrangement.
2. Only four negative charges rather than six,
which decreases electrostatic interactions
tetrahedral crystal field:
the splitting observed in a tetrahedral crystal field is
opposite of splitting in octahedral complex.
With square planar splittings,
energy level for the x2-y2 orbital is very high
so this is an especially good geometry
for d8 complexes, e.g. Pt(II), Ni(II), Pd(II), Au(III)
Factors That Affect the Magnitude of Do
Charge on the metal ion
Increasing charge on a metal ion has 2 effects:
1. Radius of metal ion decreases
2. Neg charged ligands are more strongly attracted
to it.
Both factors decrease metal-ligand distance,
which causes (-) charged ligands
to interact more strongly with the d orbitals.
magnitude of Do increases
as charge on metal ion increases
Factors That Affect the Magnitude of Do
Principal quantum # of the metal
For a series of complexes of metals
from same group in periodic table
with same charge and same ligands:
magnitude of Do increases with
increasing quantum #:
Factors That Affect the Magnitude of Do
Principal quantum # of the metal
Do (3d) << Do (4d) < Do (5d)
Increase in Do w/ increasing principal quantum #
is due to: larger radius of valence orbitals
going down a column.
Repulsive ligand-ligand interactions
are important for smaller metal ions,
which results in shorter M–L distances
and stronger d-orbital-ligand interactions
Colors of Transition-Metal Complexes
Striking colors exhibited by transition-metal complexes
are caused by the excitation of an e- from
a lower-lying d orbital to
a higher-energy d orbital,
which is called a d-d transition
For a photon to affect
the d-d transition,
its E must be = to the
difference in E btwn
the two d orbitals,
which depends on the
magnitude of Do
which depends on the
structure of the
complex.
The energy of a photon of light is inversely
proportional to its wavelength
E = hc = hu
l
Colors of Transition Metal Complexes
CFT helps explain diff colors observed
for complexes
A transition metal complex
absorbs specific l of light
Color observed is complimentary
to what was absorbed
Observed color is due to transmitted or reflected light
that is complementary in color to light that is absorbed
Rubies & Emeralds both contain
Cr3+ impurities
in octahedral 6-oxide environment.
Host lattice causes differences in
distances of d-orbital-to-ligand lengths.
Applications:
Chelating Agents:
EDTA used to treat victims of heavy metal
poisoning
Chemical Analysis:
Dimethylglyoxime turns red in presence of Ni(II)
and yellow in the presence of Pd(II).
Thiocyanate  blood-red in presence of Fe(III)
and blue in the presence of Co(II).
Applications:
Coloring Agents:
e.g. Iron blue - found in ink, paint, cosmetics
(eye shadow) and blueprints.
mixture of hexacyano complexes of
Fe(II) & Fe(III).
Drug Therapy: Cisplatin is a cancer
chemotherapeutic agent - the two chlorine
ligands get replaced by donor atoms on the
DNA double helix.
Biomolecules:
Hemoglobin and cytochrome c contain Fe-heme
complexes.
Chlorophyll contains a Mg-porphyrin complex.