CHM 003 Lecture note Transition Metal.pdf

LECTURE NOTE
CHM 003: INORGANIC CHEMISTRY -
NOTE 1 (TRANSITION ELEMENTS)

SYNOPSIS/CONTENT
• First row transition elements
• Definition of transition elements;
• Electronic configuration
• Atomic radius
• Ionization potential
• Variable oxidation state
• Formation of metal complexes
• Introduction to Coordination Chemistry
• Definition of metal complex and ligands
• Types of ligands
• Bonding in Metals complexes (Chain theory & its limitations,
Werner’s theory)
• Valence Bond theory and hybridization concept
• Study of structure and magnetic properties of octahedral and
tetrahedral complexes
• Nomenclature of coordination compounds

DEFINITION OF TRANSITION ELEMENT
• The elements in groups 3–12 as shown in the schematic
periodic table below are defined as the so-called d-block
metals.
• The term transition metal is also often used to describe this
group of elements.
• However, the IUPAC (International Union of Pure and Applied
Chemistry) defines transition metals as elements with an
incomplete d subshell or elements that can form a cation with
an incomplete d subshell.
• Therefore, the group 12
metals zinc (Zn),
cadmium (Cd) and
mercury (Hg) are not
typically classified as
transition metals.

FIRST ROW TRANSITION ELEMENTS
• Three important points to note:
• each group of d-block metals consists of three
members and is called a triad;
• metals of the second and third rows are sometimes
called the heavier d-block metals;
• Ru, Os, Rh, Ir, Pd and Pt are collectively known as the
platinum-group metals.
• Our focus will be on the ﬁrst role transition
elements.

ELECTRONIC
CONFIGURATION
• The observed ground state electronic
configurations of the first row d-block metal
atoms almost correspond to the progressive filling
of the 3d, atomic orbitals respectively.
• However, there are minor deviations from this
pattern, e.g. in the first row, the ground state
of chromium (atomic number 24) is [Ar]4s13d5 rather
than [Ar]4s23d4.
• Same thing is witnessed in Cu (atomic number 29)
is [Ar]4s13d10 rather than [Ar]4s23d9.
• A reason for this deviation is fairly complicated
but related to the energy difference between the
3d and 4s atomic orbitals and the stability of 3d
when it is half or fully filled.

IONIC FORM
• d-Block metals can show several
oxidation states as their valence
electrons can be present in more than
one atomic orbital.
• M2+ and M3+ ions of the ﬁrst-row d-
block metals follow the general
formula [Ar]3dn which implies that the
electron loss is at the 4s orbital.
• Thus, the comparative chemistry of
these metals is largely concerned with
the consequences of the successive
ﬁlling of the 3d orbitals.

CLASS WORK
• Are scandium and zinc ions
transition elements?
• Justify your answer.

ANSWER
• NO!
• The only stable ions are Sc3+ and Zn2+
• Sc = 1s22s22p63s23p64s23d1
• Thus Sc3+ = 1s22s22p63s23p6
• Sc3+ has no electron in the d-orbital,
• Zn = 1s22s22p63s23p64s23d10
• Thus Zn2+ = 1s22s22p63s23p64s03d10
• Zn2+ has a ﬁlled d-orbital

QUESTION
1. The electronic configuration of chromium is
(a) [Ar] 3d44s2
(b) [Ar] 3d5 4s1
(c) [Ar] 3d6 4s0
(d) [Ar] 3d34s2
2. Give the electronic conﬁguration of the following atoms
and ions:
(i) V (ii) Cu (iii) Zn (iv) Co2+ (v) Cr3+ (vi) Fe3+

PROPERTIES
• Nearly all the d-block metals are:
• They possess metallic lustre, high density,
• hard, ductile, high tensile strength, brittle and
malleable,
• high electrical; thermal conductivities; high m.p. and
b.p.
• They exhibit all the three types of structures: face
centred cubic (fcc), hexagonal closed packed (hcp)
and body centred cubic (bcc).
• d-Block metals tend to readily form complexes with a
characteristic colour if their ground-state electronic
conﬁguration is different from d0 or d10.

• Complex formation is often characterised by a colour
change.
• Exhibit interesting magnetic properties
• Paramagnetism
• Paramagnetism is a phenomenon that is often
observed for d-block metal compounds.
• This is a result of the presence of unpaired electrons
and can be investigated using electron paramagnetic
resonance (EPR) spectroscopy.
• As a result, abnormalities in the NMR spectra can be
observed such as broadening of the signals or unusual
chemical shifts.
Paramagnetism is defined as the phenomenon whereby some materials show magnetic
properties only once they are exposed to a magnetic field. Outside this magnetic field, no
magnetic properties are seen. This is in contrast to ferromagnets, which show magnetic
properties independent of the environment.

PROPERTIES OF THE TRANSITION METALS

QUESTION
1.The melting/boiling point of ﬁrst row
transition metals across the period
(a) Decrease across the period
(b) Decrease up to d4, then increase
(c) Remains constant
(d) Increase up to d4, then decrease
2. Which of these ions is not coloured?
Explain why?
(A) [Co(SCN)4]2-
(B) [Fe(H2O)5SCN]2+
(C) [Al(OH)4]2-
(D) Ni(DMG)2

ATOMIC RADIUS
• The atomic radii for transition metals are
smaller than their corresponding s-block
elements.
• The atomic radii of transition elements of the
ﬁrst row decrease with increase in atomic
number but this decrease becomes small
after midway.
• The atomic
radius gradually
decreases from
Sc to Cr but from
Cr to Cu, it is
nearly the same.

• The ﬁgure below illustrates that values of rmetal:
• show little variation across a given row of the d-block;
• are greater for second and third row metals than for ﬁrst row metals;
• are similar for the second and third row metals in a given triad.
• This last observation is due to the so-called lanthanoid
contraction: the steady decrease in size along the 14 lanthanoid
metals between La and Hf

• The decrease in atomic radii in each series, in the
beginning, is due to an increase in nuclear charge
from member to member which tends to pull the ns
electrons inward, i.e., it tends to reduce the size.
• At the same time, the addition of extra electrons to
(n - 1) d-orbitals also provides the screening effect.
• As the number of d electrons increases, the
screening effect increases.
• Thus, there are operating two effects namely
• screening effect and
• nuclear charge effect which oppose each other.
• In the midway onwards of the series both these
effects become nearly equal and thus, there is no
change in atomic radii in spite of the fact that
atomic number increases gradually.

VARIABLE OXIDATION STATES
• Transition elements (except Sc and Zn) show
variable oxidation states
• Arises from the similar energies required for
removal of 4s and 3d electrons
• One or more d-orbital electrons may be used in
addition to the 4s orbitals electrons in bond
formation.

VARIABLE OXIDATION STATES
• Oxidation states of the d-block metals; the most
important and stable states are marked in blue.
• Tabulation of zero oxidation states refers to their
appearance in compounds of the metal.
• In organometallic compounds, oxidation states of less
than zero are encountered .
• An oxidation state enclosed in [ ] is rare.

• Maximum rises across row to manganese
• Maximum falls as the energy required to
remove more electrons becomes very high
• All (except scandium) have an M2+ ion
• Stability of +2 state increases across the row
• Due to increase in the 3rd Ionisation Energy

MAGNETIC PROPERTIES
• Two types of magnetic behaviour are found in substances
• diamagnetism and
• paramagnetism.
• Paramagnetic substances are attracted by the magnetic
ﬁeld and weigh more while
• Diamagnetic substances are slightly repelled by the
magnetic ﬁeld and weight less.
• As the transition metal ions generally contain one or more
unpaired electrons in them and hence their complexes are
generally paramagnetic.
• Paramagnetic character increases with increase in
number of unpaired electrons.

PARAMAGNETISM
• Paramagnetism is expressed in terms of magnetic
moment.
• The more the magnetic moment the more will be
the paramagnetic character.
• The attraction of magnetic lines of force by a
compound
• Paramagnetism occurs due to unpaired electrons
• When the electrons are all paired and there is no
attraction of magnetic lines of force, the compound
is said to be DIAMAGNETIC

PARAMAGNETISM (CONTD)
• Ferromagnetism is a phenomenon whereby certain
materials form magnets a or are strongly attracted
to magnets
• Common examples are Fe, Co and Ni
The more the number of unpaired electrons, the more
the magnetic property of the transition metal

QUESTION
• The paramagnetism of transition element
compounds is due to:
(a) Paired electrons spinning in opposite direction
(b) Unpaired electrons in the d orbitals
(c) Shared valence electrons
(d) Unshared valence electrons

IONISATION ENERGIES OF TRANSITION METALS
• The ﬁrst and second ionization energies increase
only slightly from scandium to zinc because the
effective nuclear charge increases only slightly
across the series
• There is an abnormal increase in Cr and Cu (second
ionisation energy)

IONISATION ENERGIES OF TRANSITION METALS (CNTD)
• There is an increase across the series for the third
and fourth I.E because it involves the removal of
electrons from the d-orbitals
• There is a big increase in Mn and Zn, and a
depression on Fe in their third I.E
• The fourth I.E is higher than the third in all cases,
but unusually higher in Fe

ATOMIC RADII
• A slight decrease in atomic radii as we move along the
period
• This is due to the shielding effect of the outermost electrons
by the d electrons makes the effective nuclear charge felt
gradually
• Once the d-orbital becomes more than half ﬁlled, electron
repulsion forces the d-subshell to gradually expand in size
and this causes the atomic radii to increase

QUESTION
1. Which of these metals/ions has the highest
number of unpaired electrons?
(a) Cu2+ (b) Zn (c) Fe3+ (d) Ti4+
2. Which of these transition metals has the smallest
atomic radius?
(a) Mn (b) Sc (c) Fe (d) Ti
3. The element with the largest third ionization energy
is
(a) Cr (b) Sc (c) V (d) Mn
4. The element with the highest melting point is
(a) Mn (b) Fe (c) Cr (d) Cu

QUESTION
5. Which of these metals/ions has the highest number of
unpaired electrons?
(a) Cu2+ (b) Zn (c) Fe3+ (d)Ti4+
6. The element with the widest range of oxidation
numbers are
(a) Elements close to the centre of each row
(b) elements at the extreme right of each row
(c) elements at the extreme left of each row
(d) elements with the highest
7. Which compound is expected to be diamagnetic?
(a) ZnCl2 (b) FeCl3 (c) CuCl2 (d) VCl3

FORMATION OF COMPLEX IONS
• Transition metal ions form complexes or coordination
compounds
• They form complexes with molecules or ions known as
Ligands
• Ligands are either negatively charged or they have an atom
with a lone pair of electrons
• Such atoms are known as donor atoms
• Examples include O, N, S, P amongst many others.
• The complex formation is often accompanied by a change in
colour and sometimes a change in the intensity of colour.
• Equation below shows the effect of adding concentrated
HCl to aqueous cobalt(II) ions.

COLOUR FORMATION
• Transition metal ions or compounds are usually coloured due
to the promotion of an electron from one d-orbital to another
by the absorption of visible light.
• It can be clearly explained as follows: in an isolated transition
metal atom, the ﬁve d orbitals are degenerate;
• However in a complex ion, i.e. when are ligand approaches,
the d-orbitals differ in energy;
• In the transition elements which have partly ﬁlled d-orbitals,
the transition of electron can take from one of the lower d-
orbitals to some higher d-orbital within the same subshell.
• The energy required for this transition falls in the visible
region.

SHAPES OF THE D ORBITALS
There are 5
different orbitals of
the d variety

• Placing ligands around a central ion causes the energies
of the d orbitals to change
• Some of the d orbitals gain energy and some lose energy
• In an octahedral complex, two (z2 and x2-y2) go higher
and three go lower
• In a tetrahedral complex, three (xy, xz and yz) go higher
and two go lower.

3d 3d
OCTAHEDRAL TETRAHEDRAL
Δ Ε ΔΕ

COLOUR AND D-D TRANSITION
• So 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 radiation.
• Degree of splitting depends on the CENTRAL METAL ION and
the LIGAND
• The energy difference between the levels affects how much
energy is absorbed when an electron is promoted.
• The amount of energy governs the colour of light absorbed.

COLOUR (CONTD)
• For most transition metals complexes, the
frequency of light absorbed is in the Visible region

COLOURED IONS
A solution of Copper(II)sulphate is blue because
red and yellow wavelengths are absorbed
white light
blue and green not
absorbed

USES OF TRANSITION METALS
1. Transition metals have catalytic
Properties e.g
-Fe in Haber process
- Ni in Hydrogenation reactions-margarine
-Rhodium and platinum in catalytic converters
- Vanadium (V) Oxide-Contact process (sulphuric acid)
2. Iron is hard, strong and abundant,
therefore it is used in construction and
engineering, but it is usually turned into
steel ﬁrst.

3. Titanium is as strong as steel but much lighter, and is
very resistant to corrosion, so alloys of titanium are used
in the aerospace industry, and in artiﬁcial joints, such as
hip ball and sockets.
Also, titanium dioxide is a brilliantly white compound used
in paints, plastics, paper and toothpaste.
4. Copper used in wires and cables because of its good
electrical conductivity, and used in plumbing because it is
unreactive with water. It is also used as rooﬁng materials.
The copper reacts slowly with gases and water in the air
to create a thin green layer of copper compounds. This
prevents the rest of the copper from reacting.

COORDINATION CHEMISTRY
• A coordination compound is formed by
joining independent molecules or ions (known
collectively as the ligands) to a central atom
or transition metal ion using coordinate
covalent bonds.
• Ligands are atoms, or ions, which possess
lone pairs of electrons.
• Ligands form co-ordinate bonds to the central
ion by donating a lone pair into vacant
orbitals on the central species
[Co(NH3)5Cl]Cl2
[Fe(en)2(NO2)2]2SO4

• Transition metals act as Lewis acids while the
ligands act as Lewis base

• Ligands
• classiﬁed according to the number of lone pairs used in
attaching itself to the central ion (transition metals)
• Examples
• monodentate = 1 e.g Cl¯, OH¯, CN¯, NH3, and H2O
• bidentate = 2 e.g H2NCH2CH2NH2 , C2O4
2-
• polydentate = 2 or more donor atoms

• A polydentate ligand
EDTA
CH2
N
CH2
CH2
C
C
CH2 N
CH2
CH2 C
C
O
O
O
O
O O
O
O
*
*
*
*
*
*

CO-ORDINATION NUMBER & SHAPE
the shape of a complex is governed by the number of ligands around the central ion
the co-ordination number gives the number of ligands around the central ion
a change of ligand can affect the co-ordination number
Co-ordination No. Shape Example(s)
6 Octahedral [Cu(H2O)6]2+
4 Tetrahedral [CuCl4]2-
Square planar [Pt(NH3)2Cl2]
2 Linear [Ag(NH ) ]+

Common Geometries of
Complexes
Linear
Coordination Number Geometry
2
Example: [Ag(NH3)2]+

Common Geometries of Complexes
4
tetrahedral
square planar
Example: [Ni(CN)4]2-
Examples: [Zn(NH3)4]2+,
[FeCl4]-
(characteristic of metal ions with 8 d e-’s)

Common Geometries of Complexes
6
octahedral
Examples: [Co(CN)6]3-,
[Fe(en)3]3+

COORDINATION NUMBER AND HYBRIDIZATION
• The coordination number of a metal ion in a
complex is the number of ligand donor atoms to
which the metal is directly bonded.
• The structure of a ligand strongly depends on the
coordination number as it determines the number
of spatial orientation possible in any given complex.

Hybridization
• sp - linear
• sp2 – trigonal planar
• sp3 - tetrahedral
• sp3d - trigonal bipyramidal
• sp3d2 -octahedral
linear trigonal planar
tetrahedral trigonal bipyramidal octahedral

COORDINATION NUMBER AND HYBRIDIZATION
Coordination hybridization bond angle Shape
no.
2 sp 180o Linear
3 sp2 120o Trigonal planar
4 sp3 109o Tetrahedral
4 dsp2 90o Square planar
6 d2sp3 90o Octahedral

NOMENCLATURE OF COORDINATION COMPOUNDS:
IUPAC RULES
➢ Rule # 1: The cation is listed ﬁrst, then a space, followed by the
anion.
➢ For example the cation in this is named
- [Co(NH3)5Cl]Cl2
before the anion (chloride), which is named last
➢ Rule # 2: When naming the cation complex:
• Any ligands that are part of the inner coordination sphere are listed
ﬁrst, followed immediately by the name of the metal with its
oxidation state in parentheses.
• Ligands are named in alphabetical order (When writing formular of a
complex, anions are placed before neutral molecules)
• Metal atom/ion is named last
• oxidation state given in Roman numerals follows in parentheses
• Use no spaces in complex name

NOMENCLATURE: IUPAC
RULES
• The names of anionic ligands
end with the suffix -o
• -ide suffix changed to -o
• -ite suffix changed to -ito
• -ate suffix changed to –ato

NOMENCLATURE:
IUPAC RULES
• Neutral ligands are referred to by the usual name for
the molecule
• Example
• Ethylenediamine (en)
• Exceptions
• water, H2O = aqua ammonia, NH3 = amine

NOMENCLATURE: IUPAC RULES
• Rule # 3: The following prefixes are used to indicate
the number of each type of ligand in the inner sphere:
• a. If the ligand is not chelating: di (2), tri (3), tetra (4), penta
(5), hexa (6), hepta (7), octa (8), nona (9), and deca (10).
• b. If the ligand is a chelate, the prefix is followed by the name
of the ligand in parentheses: bis (2), tris (3), tetrakis (4),
pentakis (5), hexakis (6), heptakis (7), octakis (8), nonakis (9),
and decakis (10).
• Rule # 4: Other prefixes, such as cis-and trans-, are used
to designate geometrical or optical isomers, whenever
this is necessary.

• Rule # 5: If the ligand is a bridge between two metals,
the symbol - is used as a prefix before the bridging
ligand.
• Rule # 6: The ligands in the inner coordination sphere are
listed in alphabetical order of the ligand’s root name
(ignoring any prefixes).
• Rule # 7: If the anion is a coordination compound, the
suffix -ate is used; some metals have special names in
this case: ferrate (Fe), stannate (Sn), plumbate (Pb),
argentate (Ag), and aurate (Au)

• For a complex that is cationic, use the
unmodiﬁed name for the central metal.
• If a complex is an anion, its name ends
with the –ate
-appended to name of the metal
• Some common metals are given latin
based names when they appear in
anionic complexes

Transition
Metal
Name if in Cationic
Complex
Name if in Anionic
Complex
Sc Scandium Scandate
Ti titanium titanate
V vanadium vanadate
Cr chromium chromate
Mn manganese manganate
Fe iron ferrate
Co cobalt cobaltate
Ni nickel nickelate
Cu Copper cuprate
Zn Zinc zincate

EXAMPLES
• The name of the following
coordination compounds are:
• (a) [Co(NH3)5Cl]Cl2 - pentaamminechlorocobalt(III)
chloride
• (b) [Fe(acac)3] - tris(acetylacetato)iron(III)
• (c) [Pt(en)2Cl2]Cl2 -
dichlorobis(ethylenediamine)platinum(IV) chloride,
• (d) [Ru(bpy)2Cl2] -

• [Co(OX)3]3-
• Trioxalatocobaltate(III) ion
• K4[Fe(CN)6]
• Potassium hexacyanoferrate(II)

• Try this
• [Co(en)2(H2O)(CN)]Cl2
• Na2[MoOCl4]
• [Cr(H2O)4Cl2]Cl
• K4(Ni(CN)4]

1. Aquacyanobis(ethylenediamine)cobalt(III)
chloride
2. Sodium tetrachlorooxomolybdate(IV)
3. Tetraaquadichlorochromium(III) chloride
4. Potassium tetracyanonickelate(0)

• Deﬁne the following terms in coordination chemistry:
• (i) complex ion
• (ii) ligand
• coordination number
• Name the following complex ions
• (i) [Cr(H2O)6]3+ (ii) K2[TiCl6] (iii) [Pt(H2O)4(C2O4)]Br
2 (iv) [Co(NH3)6]Cl3 (v) [Co(H2NCH2CH2NH2)2Cl2]Cl
• State the oxidation number and the coordination number
of the metal in each of these complexes above

WRITING OF FORMULARS
• Tetraaminecopper(II) ion
Step 1: (NH3)4
Step2: Cu(NH3)4
Step 3: calculate the charge on the complex
i.e 2+0 =2
• [Cu(NH3)4]2+

WRITING OF FORMULARS (CLASSWORK)
• Triamminechlorodinitroplatinum(IV) ion
• Sodium hexanitrocobaltate(III)
• Hexaaquairon(III) ion
• Tetraamminebromochlorochromium(III)ion
• Trioxalatochromate(III) ion

ANSWERS
• [PtCl(NO2)2(NH3)3]+
• Na3[Co(NO2)6]
• [Fe(H2O)6]3+

CLASSWORK I
• Write the formulas for the following coordination
compounds and predict possible
• shape(s), hybridization(s) and bond angle(s) of
each:
• (i) Tetraamminediaquacobalt(III) chloride
• (ii) Potassium tetracyanonickelate(II)
• (iii) Tris(ethylenediamine)chromium(III) chloride
• (iv) Amminebromochloronitritoplatinate(II) ion
• (v) Dichlorobis(ethylenediamine)platinum(IV)
nitrate
• (vi) Iron(III) hexacyanoferrate(II)

CLASSWORK II
• 2. Write the IUPAC names of the following
coordination compounds and predict
• possible shape(s), hybridization(s) and bond
angle(s) of each::
• (i) [Co(NH3)6]Cl3
• (ii) [Co(NH3)5Cl]Cl2
• (iii) K3[Fe(CN)6]
• (iv) K3[Fe(C2O4)3]
• (v) K2[PdCl4]
• (vi) [Pt(NH3)2Cl(NH2CH3)]Cl

BONDING IN METAL COMPLEXES
• There are three theories of metal to ligand bonding
in complexes
• 1. Werner’s Theory
• 2. Valence bond Theory
• 3. Crystal ﬁeld Theory
• 4. Molecular orbital Theory

WERNER’S THEORY
• A famous scientist – Alfred Werner in 1823 put
forward his theory of coordination compounds
which describes the formation and structure of
complex compounds
• Due to this theory he is awarded by Nobel prize and
he is also called the ‘Father of Coordination
Chemistry’.

POSTULATES OF WERNER’S THEORY
• The important postulates of Werner’s theory are as
follows:
• In coordination compounds, the central metal or metal
atoms exhibit two types of valency -The primary valency
and the secondary valency.
• The primary valency corresponds to oxidation state and
the secondary valency corresponds to coordination
number.
• Every metal atom has a ﬁxed number of secondary
valencies, For Example: It has ﬁxed coordinate number.

VALENCE BOND THEORY (VBT)
• Valence bond theory is a synthesis of early
understandings of how covalent bonds form.
• It explains the reason for the variation in coordination
number based on the number of hybridized orbitals of
the metal used in bonding.
• Orbital hybridization occurs when bonding orbitals share
the characteristics of several types of orbitals.
• VBT cannot explain fully the concept of colour and
magnetic properties of complexes

CRYSTAL FIELD THEORY (CFT)
• This theory views the attractions between a central atom
or ion and its ligands being largely electrostatic
• The model focuses on the energies of the d orbitals
• In its simplest form, the theory assumes that the ligands
are negative point charged and the metal-ion bonding is
entirely ionic.

• Electrostatic Interactions
• (+) metal ion attracted to (-) ligands (anion or dipole)
• provides stability
• lone pair e-’s on ligands repulsed by e-’s in metal d
orbitals
• interaction called crystal ﬁeld
• influences d orbital energies
• not all d orbitals influenced the same way

CFT
• The crystal ﬁeld model allows us to account for the
differences in the magnetic properties of
complexes. For example, [Co(NH3)6]3+ and [CoF6]3-
• [Co(NH3)6]3+ is known to be diamagnetic while
[CoF6]3- is know to be paramgnetic because it has
four unpaired electrons

CHM 003 Lecture note Transition Metal.pdf

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