2. What is a transition element ?
• Transition element is an element having a d orbital
that is partially filled with electrons, or an element that
has the ability to form stable cation(s) with an
incompletely filled d orbital.
• The elements lying in the middle of the Periodic Table
between group 2 and group 13 are known as the d-block
elements.
3. Introduction
• These elements are known as transition elements
because they exhibit transitional behaviour
between s-block and p-block.
• Depending upon the subshell (3d, 4d, 5d) involved,
transition elements are mainly classified into three
series:
First transition series or 3d series
Second transition series or 4d series
Third transition series or 5d series
4. Electronic configuration
• The external electronic configuration is consistent.
• There is a gradual filling of 3d orbitals across the series starting
from scandium.
• However, this filling is not regular, since, at chromium and
copper, the population of 3d orbitals increases by acquiring an
electron from the 4s shell.
• At chromium, both the 3d and 4s orbitals are occupied, but
neither of the orbitals is completely filled.
• This indicates that the energies of the 3d and 4s orbitals are
relatively close for atoms in this row.
5. Electronic configuration
• The electronic configurations of first, second, and third series elements are as
follows:
First series: 1s22s2p63s2p6d1-104s2
Second series: 1s22s2p63s2p6d1-104s2p6d1-105s2
Third series: 1s22s2p63s2p6d1-104s2p6d1-10 5s2p6d1-106s2
• These three series of elements depend on the n-1 d orbital that is being filled.
An orbital of lower energy is filled first.
• Therefore, 4s orbital with lesser energy is filled first to its full degree.
• After 4s, the 3d orbital with higher energy is filled. The precisely, half-filled and
totally filled d-orbitals are exceptionally stable.
6. Transition metal electron
configuration
Anomalies do occur in all
the series, which can be
explained from the
following considerations.
The energy gap
between the ns and (n-
1) d orbitals
Pairing energy for the
electrons in s-orbital
Stability of half-filled
orbitals to the partially
filled orbitals
7. Metallic character…
• All the transition metal elements are metallic in nature.
• Transition metals crystallize in all the three face centred
cubic (fcc), hexagonal close packed (hcp) and body
centred cubic (bcc) crystals.
• Due to their greater effective nuclear charge and the large
number of valence electrons, the metallic bond is quite
strong and hence they are hard, posses high densities and
high enthalpies of atomization.
8. Metallic character
• Transition metals are hard, malleable and
ductile due to presence of strong metallic
bonds.
• Along with metallic bonding, transition metals
also show covalent bonding due to presence
of unfilled d-orbitals.
• As transition elements are metals so they good
conductors of heat and electricity
9. Atomic radius/metallic
radius
• The atomic radii of
the elements
decrease from left to
right across a row in
the transition series.
• This is because of
the poor screening by
the d electrons due to
which, the nuclear
charge attracts all of
the electrons more
strongly, hence a
contraction in size
occurs.
• The atomic radii for the elements
from Cr to Cu are very close to
one another.
10. Atomic radius…
• This closeness in atomic radii is due to the shielding of outer 4s
electrons by 3d electrons from the inward pulls of nucleus.
• As a result of these two opposing effects, the atomic radii do not alter
much on moving from Cr to Cu.
• The elements in the first group in the d-block show the excepted
increase (due to the addition of extra shell) in size Sc → Y → La.
• However in the subsequent groups there is an increase between first
and second members, but hardly any increase between second and
third elements. This is due to lanthanide contraction
12. Density
• The trend in density will be reverse of atomic radii, i.e. density increase
remains almost the same and then decreases along the period.
• Down the group density of 4d series is larger than 3d. Due to lanthanide
contraction and a larger decrease in atomic radii and hence, the volume
density of 5d series transition elements are double than 4d series.
• In the 3d series, scandium has the lowest density and copper highest density.
Osmium (d=22.57g cm-3) and Iridium (d=22.61g cm-3) of 5d series have the
highest density among all d block elements.
14. Transition elements as noble metals
• The ionization energies of elements increase very slowly across a given row.
• From the left of 3d series to the right corner 5d transition elements, density,
electronegativity, electrical and thermal conductivities increase, while
enthalpies of hydration of the metal cations decrease in magnitude.
• This indicates that the transition metals become steadily less reactive and
more “noble” in character.
• The relatively high ionization energies, increasing electronegativity, and
decreasing low enthalpies of hydration make metals (Pt, Au) in the lower
right corner of the d block as unreactive that they are often called the “noble
metals.”
15. Melting and Boiling points
• Melting and boiling points of the transition elements are
generally high.
• In general, the melting and boiling points increase from the
beginning towards the middle of the transition series.
• This is followed by a decrease, due to electrons pairing toward
the end of the series.
Reason Behind the High Melting/Boiling Points of Transition
Elements.
• The presence of unpaired electrons leads to the formation of
metal-metal covalent bonds along with the metallic bonds.
• These strong bonds attribute high melting and boiling points to
the elements.
16. Reason Behind the High Melting/Boiling Points of
Transition Elements…
• The presence of a partially filled d-orbital enables the transition elements to have a
greater number of unpaired electrons, which in turn increases their ability to form
covalent bonds along with metallic bonds.
• For example, the elements with the greatest number of unpaired electrons
(chromium, molybdenum, and tungsten) have the greatest melting and boiling
points in their respective rows.
• On the other hand, metals such as zinc and mercury do not hold any unpaired
electrons and hence have relatively low boiling and melting points.
17. Ionization energy
• The first ionization
energy of transition
elements are higher
than those of s-block
elements but lower
than p-block
elements.
• Ionization energy
increases gradually as
we move from left to
right but this increase
is not appreciable.
• The increase in
ionization energy is
due to increase in
nuclear charge, the
effect of increase in
nuclear charge is
partly balanced by the
increase in screening
effect.
18. Ionization Energy…
In transition elements, on moving along the period, the addition of the extra
electron in the (n-1) d level takes place.
This electron provides a screening effect and shields the outer ns electrons
from the nucleus pull.
Because of this shielding effect of d electrons, the effect of nuclear charge
(effective nuclear charge) on outer ns electrons is somewhat less than the
actual nuclear charge.
•
Thus the effects of the increasing nuclear charge and the shielding effect
created due to the expansion of (n-1)d orbital oppose each other.
19. Ionization Energy…
Consequently, the increase in ionization energy along the period of d-block
elements is very small.
On account of these counter effects, the ionization potentials increase
rather slowly on moving in a period of the first transition series.
• Ionization energy does not decrease going down though (contrasting main
group elements).
This is due to the insufficient shielding from f orbital electrons while still
increasing the proton number
20.
21. Variable oxidation state
• All the transition elements, apart from the first and the last, display
various oxidation states.
• Reason of variable oxidation state in d-block elements is that there is
a very small energy difference in between (n-1)d and ns orbitals.
• As a result, electrons of (n-1)d orbitals as well as ns-orbitals take part
in bond formation.
• Prior to starting of transition elements, 4s orbital has lower energy
than 3d. Consequently, 4s orbital is filled prior to 3d orbitals.
22. Variable oxidation state…
• But once a transition element is formed by putting
electron(s) in the 3d orbitals, the removal of 4s
electron(s) requires lesser energy than the 3d electron
as the electronic repulsions amongst 3d and 4s
electrons raises the energy of the 4s electrons.
• Consequently, in going from the elemental form to ionic
form, it is the 4s electrons which are removed prior to 3d
electrons.
• The table below shows common oxidation states of
elements in the first transition series.
23.
24. Trends in oxidation states
a) The minimum Oxidation state of 1 is shown by Cr, Cu, Ag, Au and Hg.
b) More stable Oxidation state increases in the order 3d ˂ 4d ˂ 5d. 3d series
elements are most stable in +2; 4d series in +2 and +4 and 5d series in +4.
Cr6+ and Mn7+ (of 3d) are not stable in their higher OS.
• Compounds containing them, CrO4
2- and MnO4
– are very reactive and strong
oxidizing agents.
• While Mo6+ and Tc7+ (of 4d) are stable in their higher OS. Compounds
containing them, MoO4
2- and TcO4
– are unreactive and stable.
• Similarly, W6+ and Re7+ (of 5d) are stable in their higher OS. Compounds
containing them, WO4
2- and ReO4
– are unreactive and stable.
25. Trends in oxidation states…
• Cations of the second and third-row transition metals in
lower oxidation states (+2 and +3) are much more easily
oxidized than the corresponding ions of the first-row
transition metals.
• For example, the most stable compounds of chromium
are those of Cr(III), but the corresponding Mo(III) and
W(III) compounds are highly reactive.
• Infact, the heavier elements in each group form stable
compounds in higher oxidation states that have no
analogues with the lightest member of the group.
26. Trends in oxidation states…
• c) Strongly oxidizing, high oxidation number elements form
compounds of oxides and fluorides and not bromides and iodides.
Vanadium form only VO4
–, CrO4
2-, MnO4
–, VF5, VCl5, VBr3, VI3 and
not VBr5, VI5. V5+ oxidizes Br– and I– to Br2 and I2 but not fluoride
because of its high electronegativity and small size.
Similarly, strongly reducing, low oxidation number elements form
bromides and iodides and not oxides and fluorides.
• d) Maximum oxidation state equal to the s and d-electrons is
exhibited by middle-order elements in each series.
• Thus, manganese in 3d series has +7, Ru in 4d and Os in 5d
possess +8 maximum oxidation state.
• e) Elements may show all the Oxidation states in between the
minimum and maximum.
27. Trends in oxidation states…
f) Elements in their lower oxidation states will be ionic and basic (TiO,
VO, CrO, MnO, TiCl2 and VCl2) in-between state amphoteric (Ti2O3,
V2O3, Mn2O3, CrO3, Cr2O3, TiCl3, VCl3 ) and higher oxidation state
covalent and acidic (V2O5, MnO3, Mn2O7, VCl4 and VOCl3 ) .
g) Lower oxidation state may get stabilized by back bonding in
complexes. Ni(CO)4, Fe(CO)5, [Ag(CN)2]–, [Ag(NH3)2]+.
Lower oxidation states in these metals are stabilized by ligands like
CO, which are pi-electron donors, whereas the higher oxidation states
are stabilized by electronegative elements like Fluorine(F) and
Oxygen(O).
Hence the high oxidation compounds of these metals are mainly
fluorides and oxides.
28. Trends in oxidation states…
h) Relative stabilities of the oxidation states depend on many
factors, like, the stability of the resulting orbital, IE,
electronegativity, enthalpy of atomization, enthalpy of
hydration, etc.
Ti4+ (3d0) is more stable than Ti3+(3d1). Mn2+ (3d5) is more stable
than Mn3+(3d4).
Ionization energies contribute to the relative stability of transition
metal compounds (ions).
For example, Ni2+ compounds are thermodynamically
more stable than Pt2+, Whereas, Pt4+ compounds are more stable
than Ni4+.
29. Trends in oxidation states…
The relative stabilities can be explained as follows:
• Thus, the ionization of Ni to Ni2+ requires lesser energy (2490 kJ mol−1) as
compared to the energy required for the production of Pt2+ (2660 kJ mol−1).
Therefore, Ni2+ compounds are thermodynamically more stable
than Pt2+ compounds.
• On the other hand, the formation of Pt4+ requires lesser energy (9360
kJ mol1) as compared to that required for the formation of Ni4+ (11290
kJ/mol). Therefore, Pt4+ compounds are more stable than Ni4+ compounds.
This is supported by the fact that [PtCl6]2+ complex ion is known, while the
corresponding ion for nickel is not known.
i) In p-block, the heavier elements prefer lower oxidation states due to what is
called the inert pair effect. But in the case of d block elements, the higher
oxidation states are more stable for heavier members in a group.
Metal (IE1+IE2) kJmol
−1
, (IE3+IE4) kJmol
−1
, Etotal, =(=IE1+IE2+
IE3+IE4) kJ mol
−1
Ni 2490 8800 11290
Pt 2660 6700 9360
30. Catalytic property
• Many transition metals and their compounds show profound catalytic
activity in many chemical and biological reactions.
• Examples: Iron in Haber’s process to make ammonia, vanadium
pentoxide in the manufacture of sulphuric acid, titanium chloride
as Ziegler-Natta catalyst in polymerization and Palladium chloride in
the conversion of ethylene to acetaldehyde are some very important
commercial catalytic processes involving d block metals.
Most transition elements act as good catalysts because of,
• The presence of vacant d-orbitals.
• The tendency to exhibit variable oxidation states.
• The tendency to form reaction intermediates with reactants.
• The presence of defects in their crystal lattices.
31. How do they catalyze reactions
They take the reaction through a path of low activation energy by:
• Providing a large surface area for absorption and allowing sufficient
time to react,
• May interact with the reactants through their empty orbitals.
• May actively interact by redox reaction through their multiple
oxidation states.
32. Coloured compounds
• Compounds of transition
elements are usually coloured
due to the promotion of an
electron from one d-orbital to
another by the absorption of
visible light.
This can be clearly explained as
follows:
• In the transition elements which
have partially filled d-orbitals,
the transition of electron to
some higher orbital within the
same subshell.
• The energy required for this
transition falls in the visible
region.
• When white light falls on these
complexes they absorb a
particular colour from the
radiation for the promotion of
electron and the remaining
colours are emitted.
• The colour of the complex is
due to this emitted light.
• A few of the transition metal
ions such as Cu+, Ag+, Sc3+
are colourless. In these ions,
the d-orbitals are either
completely filled or empty.
33. Formation of coloured ions
• a) Colour of the ions varies with their oxidation state. The Cr6+ as
in potassium dichromate is yellow in colour, whereas
Cr3+ and Cr2+ are generally green and blue respectively.
• b) Colour of the compound is dependent on the complexing or
coordinating group also.
• For example, Cu2+ shows light blue colour in the presence of water
as ligand but the deep blue colour in the presence of ammonai as the
ligand.
• c) Transition metal ions which have:
• Completely filled d-orbitals having no vacant d-orbitals for excitation
of electrons are colourless. Cu+(3d10), Zn2+(3d10), Cd2+(4d10)
Hg2+(5d10), and Zn, Cd, Hg are colourless.
• Transition metal ions which have completely empty d-orbitals without
d-electrons are also colourless. Sc3++(3d0), and Ti4++(3d0), ions are
colorless.
34. Complex formation
• Transition metal complexes are species that contain groups with lone pair(s)
of electrons arranged around a central metal ion.
• The species that attach to the transition metal atom are called ligands.
• Ligands may be molecules or ions that contain lone pairs of electrons
• Common examples are: water, hydroxide, halogens, cyanide and ammonia.
• Transition metal ions form variety of complexes due to the following reasons:
i) Small size and high nuclear charge
ii) Availability of vacant d-orbital of suitable energy, which can accept lone
pair(s) of electrons donated by ligands.
Examples of transition metal complexes
[Ag(NH3)2]+
[Fe(H2O)6]3+
[CuCl4]3-
35. Paramagnetism
• Transition elements show paramagnetism due to presence of one or more
unpaired electrons in them. What is diamagnetism?
• Paramagnetism is when a substance is weakly attracted to a magnetic field.
• Paramagnetic substances are attracted weakly by magnetic field and weigh
more while diamagnetic substances are slightly repelled by magnetic field
and weigh less.
• As the transition metal ions generally contain one or more unpaired
electrons in them, they are generally paramagnetic.
• Paramagnetic character increases with increase in number of unpaired
electrons.
• Paramagnetism is expressed in terms of magnetic moment as
μ = √[n(n+2)] BM
Where n= number of unpaired electrons, B=magnetic field strength and is the
magnetic moment.
36. Determination of Paramagnetsim
o The instrument consists of a long,
cylindrical sample suspended
from a balance, partially entering
between the poles of a strong
magnet.
o The balance measures the
apparent change in the mass of
the sample as it is repelled or
attracted by the region of high
magnetic field between the poles.
o In a practical device the whole
assembly of balance and magnet
is enclosed in a glass box to
ensure that the weight
measurement is not affected by
air currents.
o The sample can also be enclosed
in a thermostat in order to make
measurements at different
temperatures
o A Gouy balance is a device
that can be used to measure
the weak diamagnetic and
paramagnetic effects of
transition metal compounds.
37. Determination of magnetic susceptibility using a Gouy
balance
• The Gouy balance measures the apparent change in the mass of the sample as
it is repelled or attracted by the region of high magnetic between the poles.
• Some commercially available balances have a port at their base for this
application. In use, a long, cylindrical sample to be tested is suspended from a
balance partially entering between the poles of a magnet.
• The sample can be in solid or liquid form, and is often placed in a cylindrical
container such as a test tube.
• Solid compounds are generally ground into a fine powder to allow for uniformity
amongst the sample.
• The sample is suspended between the magnetic poles through an attached
thread or string.
• The experimental procedure requires two separate reading to be performed. An
initial balance reading is performed on the sample of interest without a magnetic
field (m a).
• A subsequent balance reading is taken with an applied magnetic field (mb).
• The difference between these two readings relates to the magnetic force on the
sample (mb – ma).
38. Determination of magnetic susceptibility
using a Gouy balance
The following mathematical equation relates the apparent change in mass
to the volume susceptibility of the sample:
Force= (mb – ma)g
= ½ (K2 - K1)A H2
mb – ma = apparent difference in mass
g = gravitational acceleration
K1 = volume susceptibility of medium, usually air and of negligible value
K2 = volume susceptibility of sample
H = applied magnetic field
A = area of the sample tube
39. Transition metal complex formation
• A complex ion (or, simply, complex) is a species formed between a central
metal ion and one or more surrounding ligands, molecules or ions that
contain at least one lone pair of electrons.
• It is formed from a metal ion and a ligand because of a Lewis acid–base
interaction.
• The positively charged metal ion acts as a Lewis acid, and the ligand, with
one or more lone pairs of electrons, acts as a Lewis base.
• Small, highly charged metal ions, such as Cu2+ or Ru3+, have the greatest
tendency to act as Lewis acids, and consequently, they have the greatest
tendency to form complex ions.
40. Transition metal complex formation
• As an example of the formation of complex ions, consider the addition of
ammonia to an aqueous solution of the hydrated Cu2+ ion [Cu(H2O)6]2+.
• Because it is a stronger base than H2O, ammonia replaces the water
molecules in the hydrated ion to form the [Cu(NH3)4(H2O)2]2+ ion.
• Formation of the [Cu(NH3)4(H2O)2]2+ complex is accompanied by a dramatic
color change.
• The solution changes from the light blue of [Cu(H2O)6]2+ to the blue-violet
characteristic of the [Cu(NH3)4(H2O)2]2+ ion.
41. Effect of formation of Complex Ions on
Solubility
Formation of complex ions can substantially increase the solubility of
sparingly soluble salts if the complex ion has a large Kf.
A complex ion is a species formed between a central metal ion and one or
more surrounding ligands, molecules or ions that contain at least one lone
pair of electrons.
Small, highly charged metal ions have the greatest tendency to act as Lewis
acids and form complex ions.
The equilibrium constant for the formation of the complex ion is the formation
constant (Kf).
The formation of a complex ion by adding a complexing agent increases the
solubility of a compound.
42. Effect of formation of Complex Ions on
Solubility and applications
Complexing agents, molecules or ions that increase the solubility of metal
salts by forming soluble metal complexes, are common components of
laundry detergents.
Long-chain carboxylic acids, the major components of soaps, form insoluble
salts with Ca2+ and Mg2+, which are present in high concentrations in “hard”
water. The precipitation of these salts produces a bathtub ring and gives a
gray tinge to clothing.
Adding a complexing agent such as pyrophosphate (O3POPO3
4−, or P2O7
4−)
or triphosphate (P3O10
5−) to detergents prevents the magnesium and calcium
salts from precipitating because the equilibrium constant for complex-ion
formation is large.
Ca2+
(aq) + O3POPO4
4-
(aq) ↔ [Ca(O3POPO3)]2−
(aq)
43. Effect of the formation of complex ions on
Solubility and applications
Another application of complexing agents is found in medicine.
Magnetic resonance imaging (MRI) can give relatively good images of soft
tissues such as internal organs.
MRI is based on the magnetic properties of the 1H nucleus of hydrogen
atoms in water, which is a major component of soft tissues.
Scientists have developed a class of metal complexes known as “MRI
contrast agents.”
Injecting an MRI contrast agent into a patient selectively affects the magnetic
properties of water in cells of normal tissues, in tumors, or in blood vessels
and allows doctors to “see” each of these separately.
One of the most important metal ions for this application is Gd3+, which with
seven unpaired electrons is highly paramagnetic. Because Gd3+(aq) is quite
toxic, it must be administered as a very stable complex that does not
dissociate in the body and can be excreted intact by the kidneys. The
complexing agents used for gadolinium are ligands such as
DTPA5− (diethylene triamine pentaacetic acid), whose fully protonated form.
44. Coordination compounds and complex ions
Coordination compound is a chemical species containing a central metal
atom surrounded by nonmetal atoms or group of atoms called ligands ,
joined to it by chemical bonds.
They are called coordination compounds because there are coordinate
covalent bondsbetween metal atoms and ligands.
Examples of coordination compounds are as follows:
• Haemoglobin
• Chlorophyll
• Dyes
• Pigments
• Vitamin B12
• Enzymes etc
Complex ion is a coordination compound with an electrical charge.
Therefore, this chemical species also contains a central metal atom and
ligands bound to it.
45.
46. Difference between Primary and Secondary
Valency in Coordination Compounds
Primary valency Secondary valency
These are ionizable These are Non-ionizable
Satisfied by charged ions Satisfied by ligands
Primary valency does not help in
the structure of complex
Secondary valency helps in structure
It can also function as a secondary
valence
It can not function as a primary
valency
47. Metal valence
• Primary valence – the oxidation state
• Secondary valence – the number of molecules or ions directly bound to
the metal center
• Also known as the coordination number
• The silver complex ion below [Ag(NH3)2]+ has a coordination number
of 2 since two amine ligands are bound to it directly
48. Terms used to describe
Coordination compounds
• Complex ion – A
transition metal with
one or more ligands
bound to it
• Ligand – a lewis base
that can form a bond
with the metal
• Coordination
compound – a
complex ion balanced
with one or more
counterions
49. Terms cont..
• Coordination Entity
A chemical compound in which the central ion or atom (or the coordination
centre) is bound to a set number of atoms, molecules or ions is called
a coordination entity. Examples of such coordination entities include
[CoCl3(NH3)3], and [Fe(CN)6]4-.
• Central Atoms and Central Ions
• As discussed earlier, the atoms and ions to which a set number of atoms,
molecules, or ions are bound are referred to as the central atoms and
the central ions.
• In coordination compounds, the central atoms or ions are typically Lewis
acid and can, therefore, act as electron-pair acceptors.
• Ligands
• The atoms, molecules, or ions that are bound to the coordination centre or
the central atom/ion are referred to as ligands.
• These ligands can either be a simple ion or molecule (such as Cl– or NH3) or
in the form of relatively large molecules, such as ethane-1,2-diamine (NH2-
CH2-CH2-NH2).
50.
51. Terms Used…
• Coordination Number
• The coordination number of the central atom in the coordination compound
refers to the total number of sigma bonds through which the ligands are
bound to the coordination centre.
• For example, in the coordination complex given by [Ni(NH3)4]2+,
the coordination number of nickel is 4.
• Coordination Sphere
• The non-ionizable part of a complex compound which consists of central
transition metal ion surrounded by neighbouring atoms or groups enclosed
in square bracket.
• The coordination centre, the ligands attached to the coordination centre, and
the net charge of the chemical compound as a whole, form the coordination
sphere when written together.
• This coordination sphere is usually accompanied by a counter ion (the
ionizable groups that attach to charged coordination complexes).
• Example: [Co(NH3)6]C/3 – coordination sphere
52. Terms used…
• Coordination Polyhedron
• The geometric shape formed by the attachment of the ligands to the
coordination centre is called the coordination polyhedron.
• Examples of such spatial arrangements in coordination compounds include
tetrahedral and square planar.
• Oxidation Number
• The oxidation number of the central atom can be calculated by finding the
charge associated with it when all the electron pairs that are donated by the
ligands are removed from it.
• For example, the oxidation number of the platinum atom in the complex
[PtCl6]2- is +4.
• Homoleptic and Heteroleptic Complex
When the coordination centre is bound to only one type of electron pair
donating ligand group, the coordination complex is called a homoleptic
complex, for example: [Cu(CN)4]3-.
When the central atom is bound to many different types of ligands, the
coordination compound in question is called a heteroleptic complex, an
example for which is [Co(NH3)4Cl2]+.
53. Terms used…
Cationic complexes: In this co-ordination sphere is a cation. Example: [Co(NH3)6]Cl3
Anionic complexes: In this co-ordination sphere is Anion. Example: K4[Fe(CH)6]
Neutral Complexes: In this co-ordination sphere is neither cation or anion.
• Mononuclear complexes: In this co-ordination sphere has single transition metal ion.
Example: K4[Fe(CN)6]
• Polynuclear complexes: More than one transition
metal ion is present. Example:
54. Types of Ligands
• Unidentate Ligands
• The ligands which only have one atom that can bind to the coordination centre are called
unidentate ligands. Ammonia (NH3) is a great example of a unidentate ligand. Some common
unidentate are Cl–, H2O etc.
• Bidentate Ligands
• Ligands which have the ability to bind to the central atom via two separate donor atoms, such
as ethane-1,2-diamine, are referred to as bidentate ligands.
• Oxalate ion is a bidentate as it can bond through two atoms to the central atom in a
coordination compound and Ethane-1, 2-diamine:
55. Types of ligands
• Polydentate Ligands
• Some ligands have many donor atoms that can bind to the coordination
centre. These ligands are often referred to as polydentate ligands.
• A great example of a polydentate ligand is the EDTA4- ion (ethylene diamine
tetraacetate ion), which can bind to the coordination centre via its four
oxygen atoms and two nitrogen atoms.
• Chelate Ligand
• When a polydentate ligand attaches itself to the same central metal atom
through two or more donor atoms, it is known as a chelate ligand. The atoms
that ligate to the metal ion are termed as the denticity of such ligands.
• Ambidentate Ligand
• Some ligands have the ability to bind to the central atom via the atoms of
two different elements.
• For example, the SCN– ion can bind to a ligand via the nitrogen atom or via
the sulfur atom. Such ligands are known as ambidentate ligands.
56.
57.
58.
59. Nomenclature of coordination compounds and
metal complexes
Rules to follow
1. The ligands are always written before the central metal ion in the naming of
complex coordination complexes.
2. When the coordination centre is bound to more than one ligand, the names
of the ligands are written in an alphabetical order which is not affected by the
numerical prefixes that must be applied to the ligands.
3. When there are many monodentate ligands present in the coordination
compound, the prefixes that give insight into the number of ligands are of the
type: di-, tri-, tetra-, and so on.
4. When there are many polydentate ligands attached to the central metal ion,
the prefixes are of the form bis-, tris-, and so on.
5. The names of the anions present in a coordination compound must end with
the letter ‘o’, which generally replaces the letter ‘e’. Therefore, the sulfate
anion must be written as ‘sulfato’ and the chloride anion must be written as
‘chlorido’.
60. Rules for naming coordination compounds and complexes…
6. The following neutral ligands are assigned specific names in coordination compounds:
NH3 (ammine), H2O (aqua or aquo), CO (carbonyl), NO (nitrosyl).
7. After the ligands are named, the name of the central metal atom is written. If the complex
has an anionic charge associated with it, the suffix ‘-ate’ is applied.
8. When writing the name of the central metallic atom in an anionic complex, priority is given
to the Latin name of the metal if it exists (with the exception of mercury).
9. The oxidation state of the central metal atom/ion must be specified with the help of roman
numerals that are enclosed in a set of parentheses.
10. If the coordination compound is accompanied by a counter ion, the cationic entity must be
written before the anionic entity.
61. Rules for writing chemical formulas from the
name of complex /coordination compound
• The complex ion is written within brackets and the
overall charge is written as a superscript on the
upper right
• Within the brackets, the metal is written first, then
the ligands are written in order of neutral species
then ionic species in alphabetical order of the
chemical symbols
• Write the cation first and the anion second if it is part
of an ionic compound
• Make sure if it is part of an ionic compound that the
net charge is zero.
65. Common geometries
• Depending on the amount of ligands coordinated to the metal
center a different geometry will result
• Geometry is a property that plays can play a crucial role in
reactivity
– Must consider the physical 3D component of a molecule in
many cases
• A complex with a high coordination number less likely to
have additional ligands coordinate
• A complex with a low coordination number more likely to
have additional ligands coordinate
66.
67. Structure and Isomerization
• Isomer – molecules with the same formula, but a different arrangement of
atoms (different structure)
• Structure determines the properties of a compound
• Structural isomers – atoms connected in different ways to each other
• Stereoisomers – atoms connected in the same way, but different spatial
arrangement
70. Structural Isomers
• Linkage isomers – contain ligands that can
coordinate to the metal through different
atoms
– How a ligand coordinate affects the name of the
ligand
72. Stereoisomers
• Geometric isomers – ligands coordinate to
metal in different spatial arrangement
• Square planar and octahedral complexes that
have two identical ligands can have cis or trans
orientation
– cis – when the identical ligands are next to each
other
– trans – when the identical ligands are opposite of
each other
73. Cis/Trans Isomers (Type of Stereoisomer)
Chloride ligands
are cis to each
other
Chloride
ligands are
trans to each
other
75. Stereoisomers
• Optical isomers –
Mirror images that are
nonsuperimposable
– Like your left and right
hand, they are mirror
images of each other,
but are not the
same/superimposable
76. Since the mirror image is
superimposable there is no
optical isomer
77. Bonding in coordination compounds
• Geometries are derived from the orbitals,
since orbitals interact in bonding
• Valence bond theory – filled orbital (ligand)
interacts with unfilled orbital (metal)
– Metal ion orbitals hybridize to produce
appropriate geometries
• This model helps to explain shape but not
properties such as color and magnetic
properties
78.
79. Crystal Field Theory
• Crystal Field Theory (CFT) – focuses on interaction
of the ligand with the metal d orbitals
• Ligands are like an electron cloud attracted to a
positively charged metal center
• When the ligands approach metal center, they
repel the electrons/potential electrons in the
unhybridized metal d orbitals
• This repulsion increases the energy of the d
orbitals
– the interaction between d orbitals differs so the
energy levels are split
• Crytal field splitting energy = Δ
– Difference between higher and lower energy levels
80.
81. d orbital splitting in an octahedral
complex
• z2 and x2-y2 d orbitals have largest interaction, so
they increase by a greater amount
86. Color in coordination compounds
• Where does color come from in coordination
compounds?
• Depending on the light absorbed at different
wavelengths, a certain color is observed
• When light is absorbed, an electron in a lower
d orbital can transition to a higher d orbital
Light photon
Electron excited to higher
energy orbital
87. Wavelength of color absorbed produces
complementary color to be observed
Red absorbed
Green observed
Violet absorbed
Yellow observed
89. Color explained!
• Depending on the difference in energy levels,
the energy absorbed to excite an electron is
changed, producing a corresponding color
observed
• Ligands dictate the energy levels based upon
interaction
• Strong field ligands - interact strongly to
increase the splitting energy more
• Weak field ligands - interact weakly increasing
the splitting energy less
90. Spectrochemical series
• Stronger field ligands increase splitting energy
in comparison to weaker field ligands
• Spectrochemical series is a list of ligands in
order of decreasing ligands strength
91. Practice
• From each pair, which complex ion will have a
larger crystal field splitting (Δ)
1. Fe(CN)6
2- or Fe(H2O)6
2+
2. Co(NO2)Cl3
- or Co(OH)Cl3
-
3. CrF6 or CrI6
4. Mn(en)2
2+ or Mn(NH3)2
2+
92. Magnetic properties
• If energy splitting is
significant enough,
placement of electrons
can be changed
• If all electrons are paired,
complex is diamagnetic
(low spin)
• If unpaired electrons are
present, complex is
paramagnetic (high spin)
