Transcript Lectures 29-31

Transition Metals
•Mercury (Hg) is the only transition metal that is not a
solid.
•The transition metals all have valence electrons in a d
subshell.
•Like other metals, transition metals form cations not
anions.
•We shall see that many transitions cations form
beautifully coloured compounds.
Ligands and Coordiantion Complexes
•Co-ordination complexes are compounds in which several ligands are
co-ordinated to a transition metal cation. A ligand is any substance
(neutral or anion) which can act as a Lewis base, donating electrons to
the transition metal cation (which acts as a Lewis acid). If the complex
has a charge, it is a complex ion.
•[Cu(OH2)6]2+ is Cu2+ with six water (“aqua ligands”)
2+
OH 2
H 2O
OH2
Cu
H2 O
OH 2
OH 2
•[Zn(CN)4]2- is Zn2+ with four cyanide (“cyano ligands”)
2NC
CN
Zn
NC
CN
The ligands around the metal do not all have to be the same!
Ligands and Coordiantion Complexes
•A very important co-ordination complex is found in hemoglobin:
This is a cartoon!
Heme (the porphyrin
in hemoglogin) has
chains branching off
the porphyrin ring.
Ligands and Coordiantion Complexes
•Classifying Ligands
•Ligands co-ordinated to a transition metal through one atom are called
monodentate ligands.
•Ligands co-ordinated to a transition metal through two atoms are
called bidentate (“two-toothed”) ligands.
•Polydentate ligands can also be called chelating ligands, or chelates
(“claws”). We saw one such ligand in the Chemistry 1000 “Hardness of
Water” lab. EDTA was able to “grip” a cation by co-ordinating to it with
six different atoms! (For clarity, individual carbon atoms are not
shown.)
3-
O
OH2
H2O
OH2
Fe
H2O
C
3+
OH2
C
O
O
C
O
O
C
O
O
O
N
O
O
Fe
O
O
OH2
C
C
O
O
monodentate ligand
1-
O
bidentate ligand
O
O
Fe
O
N
O
O
polydentate ligand
Ligands and Coordiantion Complexes
•The number of atoms attached to the transition metal is referred to as the
co-ordination number. It doesn’t matter whether these atoms come
from the same molecule/ion or from several different ones. Go back and
assign a co-ordination number to each complex ion on the previous three
pages.
•Co-ordination complexes can be charged or neutral. To make a neutral
precipitate, charged co-ordination complexes (complex ions) need one or
more counterions to balance the charge. This gives a complex salt.
•In the CHEM 2000 lab, you will make the bright green complex salt,
K3[Fe(C2O4)3].3H2O containing Fe3+. Break this formula into a complex
ion, counterion and water of hydration, clearly indicating each ion’s
charge. Identify the ligands and their charge.
3-
O
C
3 x K+ counter ions
+
O
C
O
O
C
O
C
O
O
+
Fe
O
O
O
C
O
C
O
3 x H2O water of
hydration. Water is not
coordinated to the metal
Ligands and Coordiantion Complexes
•Some co-ordination complexes and complex salts contain extra water
molecules which were trapped during crystallization. These complexes are
also hydrates. Water of hydration can be removed by heating a complex
salt in a dry oven.
•If 5.00 grams of K3[Fe(C2O4)3].3H2O is heated until all of the water has
evaporated, what mass of solid will remain?
MW of K3[Fe(C2O4)3].3H2O = 491.24274 g/mol
MW of K3[Fe(C2O4)3] = 437.1969 g/mol
Calculate the number of moles of K3[Fe(C2O4)3].3H2O (which is also the moles of
K3[Fe(C2O4)3]). Multiply the number of moles by the mass of K3[Fe(C2O4)3] to get the mass of
K3[Fe(C2O4)3] remaining.
•A co-ordination complex must contain a transition metal cation and several
ligands. It may also have counterion(s) (to balance charge) and/or extra
water molecules. When naming a co-ordination complex or complex salt,
look for these components.
Naming Complex Salts
•The first step in naming a complex salt is to identify the complex ion. To
name the complex ion:
•List the ligands using prefixes to indicate the number of each type of
ligand. Use alphabetical order if there are multiple ligands.
•For ligands with simple names (e.g. chloro, hydroxo), use di, tri,
tetra, penta, hexa, etc.
•For ligands with complicated names (e.g. oxalato), use bis, tris,
and tetrakis.
•Name the transition metal. If the complex ion is an anion, use the
metal’s Latin name and change the suffix to ‘ate’
•List the metal’s oxidation state using Roman numerals.
•Once you have named the complex ion, name the complex salt like any
other ionic compound: cation then anion then hydration.
e.g. K3[Fe(C2O4)3].3H2O
potassium trisoxalatoferrate(III) trihydrate
(cation)
(complex anion) (hydration)
Naming Complex Salts (Ligand Names)
Anions
Formula
Name
fluoride
:F-
fluoro
chloride
:Cl-
chloro
bromide
:Br-
bromo
iodide
:I-
iodo
cyanide
:CN-
cyano
oxide
:O2-
oxo
hydroxide
:OH-
hydroxo
carbonate
oxalate
[
[
O
:OCO:
OO
:OCCO:
]
]
2-
2-
carbonato
oxalato
Neutral Molecules
Formula
Name
carbon monoxide
:CO
carbonyl
water
:OH2
aqua
ammonia
:NH3
ammine
NH2CH2CH2NH2
ethylenediamine
ethylenediamine (“en”)
Naming Complex Salts (Latin Names)
When cobalt is in a complex anion, it is cobaltate. Similarly, zinc is zincate
and chromium is chromate. The elements below have names that are not
directly derived from the english name for the element.
Element
Symbol
Latin Name
Name in Anionic Complex
copper
Cu
cuprum
cuprate
gold
Au
aurum
aurate
iron
Fe
ferrum
ferrate
silver
Ag
argentum
argentate
Naming Complex Salts
•Name the following complex salts. Note that complex ions are typically
written inside square brackets.
•[Ni(OH2)6] CO3
Hexaaquanickel(II) carbonate
•[Cu(NH3)4] SO4 · H2O
Tetraaminecopper(II) sulfate monohydrate
•[CoCl3(NH3)3]
Triamminetrichlorocobalt(III)
•[Co(NH3)6] [Cr(CN)6]
Hexaamminecobalt(III) hexacyanochromate(III)
Naming Complex Salts
•Note that there is a difference between water as a ligand and “water of
crystallization”. The bright blue crystals commonly referred to as
CuSO4·5H2O are really [Cu(OH2)4]SO4·H2O. Give the name corresponding
to each of these two formulas.
•CuSO4·5H2O = Copper(II) sulfate pentahydrate
•[Cu(OH2)4]SO4·H2O = Tetraaquacopper(II) sulfate monohydrate
•The only way to determine this information is by experiment, but you
should recognize that, in many hydrated salts, at least some of the water
molecules serve as ligands.
Why are Transition Metals Special?
•We have seen that main group metals are somewhat limited in what
oxidation states they can adopt. Many transition metals, on the other
hand, can take on a wide variety of different oxidation states. This
distribution is not entirely random, as show in the graph below (common
oxidation states in dark red):
•Note that the elements in the middle can exist in a wider variety of
oxidation states than those on either end of the d-block.
Why are Transition Metals Special?
•Compared to s and p electrons, d electrons can be added or removed
relatively easily.
•The electron configuration of neutral vanadium is:
[Ar]4s23d3
•The first two electrons removed will be those in the 4s orbital. After that,
electrons are removed from the 3d orbitals giving three stable oxidation
states:
•vanadium(III) [Ar]3d2
•vanadium(IV)
[Ar]3d1
•vanadium(V)
[Ar]
Electronic Structure and Colour
One of the more fun consequences of these partially filled d subshells is
that the co-ordination complexes of transition metals are often brightly
coloured. The flasks below contain aqueous solutions of several nitrate
salts. Note that, since all nitrates are water-soluble, these solutions
contain aqua complexes of the transition metal cation.
Fe3+
Co2+
Ni2+
Cu2+
Zn2+
Electronic Structure and Colour
•Why is the Zn2+ complex the only colourless one?
•Consider the electron configurations of the five cations:
•Fe3+
[Ar]4s23d6
[Ar]3d5
Co2+
[Ar]4s23d7
[Ar]3d7
•Ni2+
[Ar]4s23d8
[Ar]3d8
•Cu2+
[Ar]4s13d10
[Ar]3d9
•Zn2+
[Ar]4s23d10
[Ar]3d10
Electronic Structure and Colour
•Where does the variety in colour come from?
•Many co-ordination complexes have octahedral geometry. This means
that two of the d orbitals on the transition metal point directly at
ligands while the other three do not:
•A simple electrostatic model, called the crystal field theory, assumes
that there will be a certain degree of electron-electron repulsion
between the electron pair a ligand donates and any electrons already in
the metal d orbitals. This repulsion is felt most strongly by electrons in
d orbitals pointing at the ligands.
Electronic Structure and Colour
•Thus, the dz2 and dx2-y2 orbitals are pushed to higher energy than the dxy,
dxz and dyz orbitals. This separation in energy is referred to as crystal
field splitting (Δo where ‘o’ is for ‘octahedral’).
Electronic Structure and Colour
In co-ordination complexes with crystal field splitting, there are two ways
to distribute d electrons. Which one is favoured depends on the size of Δo.
•The high spin distribution maximizes the alignment of spin of the d
electrons. It is favoured when Δo is small (when the metal is bonded to
weak field ligands). Why?
•The low spin distribution puts electrons in the lowest energy orbitals
first. It is favoured when Δo is large (when the metal is bonded to
strong field ligands). Why?
strong field
CN- > en > NH3 > EDTA4- > H2O > ox2- > OH- > F- > Cl- > Br- > I-
weak field
Electronic Structure and Colour
•How does this make for coloured solutions?
•Recall that photons are emitted when electrons drop from a higher
energy orbital to a lower energy orbital. (see Atomic Line Spectra)
Similarly, the electrons get to the higher energy orbital by absorbing
photons of light.
•Electrons in the lower energy d orbitals can absorb photons and be
excited into the higher energy d orbitals. Since Δo corresponds to the
energy of light in the visible region (and there is more than one way to
absorb a photon), some wavelengths of visible light are absorbed. The
wavelengths that are not absorbed give the colour of solution.
•+To absorb coloured light, the transition metal needs to have
electrons in at least one of the low-energy d orbitals and an empty
space in at least one of the high-energy d orbitals. Which of these two
requirements does Zn2+ lack (making it colourless)?
Electronic Structure and Colour
Electronic Structure and Colour
•Note that different ligands provide different amounts of crystal field
splitting. Fe(OH2)63+ and Fe(C2O4)33- are both complexes of Fe3+ but
Fe(OH2)63+ is extremely pale purple (frequently appearing colourless)
while Fe(C2O4)33- is green.
•What colour of light is each compound most likely absorbing?
•Which of these two ligands is splitting the d orbitals of Fe3+ more?
(i.e. which complex has a larger Δo)
G
Y
C
B
R
M
Isomers
•Even a very small change in the structure of a complex ion can change
its colour drastically.
•Draw two different Lewis structures for [CoCl2(NH3)4]+.
+
+
NH 3
H 3N
NH 3
Cl
Co
H3 N
Cl
NH 3
purple
H 3N
Cl
Co
Cl
NH 3
NH 3
green
•One of these compounds is purple while one is green! The purple
one is referred to as cis-[CoCl2(NH3)4]+ while the green one is trans[CoCl2(NH3)4]+
•These compounds are referred
to as isomers. They have the
same molecular formula but
one cannot be superimposed
on the other, no matter how
they are rotated.
Isomers
Draw two isomers of diamminedichloroplatinum(II), a square planar
complex.
H3 N
H3 N
Cl
Pt
H3 N
Cl
Pt
Cl
Cl
NH 3
The cis isomer is an anti-cancer drug while the trans isomer is toxic!
Draw two isomers of [CoCl3(NH3)3], an octahedral complex.
NH 3
NH 3
Cl
Cl
H 3N
Co
Co
H3N
Cl
Cl
NH 3
meridional (mer)
H3 N
Cl
Cl
facial (fac)