Lecture principles of inorganic chemistry

2P32 Winter Term 2015-16
Dr. M. Pilkington

Principles of Inorganic Chemistry

Lecture 1 – Recapping Important Concepts

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Inorganic Chemistry and the Periodic Table
Bonding Models
Shapes of Molecules - Lewis Structures
Valence bond theory: cases of NH3 H2O and BF3
Lewis Acids and Bases
σ and π bonds in CH2=CH2
The Shapes of Molecules – Relationship between Lewis
Structure, VSEPR theory and VBT.

Assignment 1 – Drawing Lewis structures and predicting the
shapes/geometries of molecules due after class Tuesday 12th January

1.
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Inorganic Chemistry and the Periodic Table
Carbon is only one element and has limited bonding modes,
oxidation states and coordination numbers.
But it does CATENATE well and forms MULTIPLE BONDS with
itself and other p-block elements especially N and O.
For the rest of the elements:
Wide range of electronegativity, oxidation states, coordination
numbers, ability to form multiple bonds and catenate etc…
How can we make sense of such wide ranging behaviors?
We have a system called the Periodic Table. The ‘Periodic Law’
1860-1870 (Mendeleev and Meyer): A periodic repetition of
physical and chemical properties occurs when the elements are
arranged in order of increasing atomic weight [number]’

1


With the development of atomic theory and spectroscopic techniques
the modern Periodic Table has evolved:

2P32 Course Outline:
Lectures 1-16
Coordination Chemistry
of transition Metal ions
Lectures 17 – 34
Descriptive Inorganic
Chemistry – Main Group
Elements.

2.

Bonding Models:

In covalent species, electrons are shared between atoms.
In an ionic species, one or more electrons are transferred between
atoms to form bonds.
Modern views of molecular structure are, based on applying wave
mechanics to molecules; such studies provide answers as to how and
why atoms combine. Two such methods are:
1. Valence Bond (VB) approach- overlap of valence orbitals on
atoms to form bonds.
2. Molecular Orbital theory (MO) of bond formation – allocates
electrons to molecular orbitals formed by the overlap (interaction)
of atomic orbitals.
Familiarity with both VB and MO concepts is necessary as it is
often the case that a given situation can be approached using one
or the other of these models.

2


3.


Shapes of Molecules
Understanding the shapes of molecules is an important step in
being able to discuss and predict chemical properties. Although
here we discuss the shapes of “simple” molecules, this topic has
also important applications in the understanding of the behavior
of much larger molecules, e.g the shape of macromolecules in

biology is often important with respect to their biochemical
function

Lewis structures – you need to be able to draw these.
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Lewis presented a simple but useful method of describing the
arrangement of valence electrons in molecules.
Lewis structures give the connectivity of an atom in a molecule,
the bond order and the number of lone pairs and these maybe
used to derive structures.
Revise your first year notes.

Test Question
 Draw the Lewis Structure of the Nitrato ion NO3-.

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How many  bonds, how many  bonds?

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What is the nitrogen-oxygen bond order?

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Are there possible resonance structures, can you draw them?

3


Bond Order
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Single bond - first order
Double bond = second order
Triple bond  third order

Bond order is a measure of the number of bonding electron pairs
between atoms. Single bonds have a bond order of 1, double bonds have
a bond order of 2 and triple bonds (the maximum number) have a bond
order of 3. A fractional bond order is possible in molecules and ions that
have resonance structures. In the example of ozone, the bond order
would be the average of a double bond and a single bond or 1.5 (3
divided by 2). As the bond order becomes larger, the bond length
becomes smaller.
Remember atoms in the 3rd period or below e.g. P, I do not always obey
the Octet rule!

4.

Valence Bond Theory

The Shape of Ammonia (NH3) – VSEPR is important here.

Lewis Structure
Lone Pair
H

N

But why isnt the NHN angle 900?

H

H

We have to consider repulsions between the lone pair and valence electrons
actual structure:
N
H

H

H-N-H angle is just slightly smaller than 109.50
The Nitrogen atom is Pyramidal

H

Ammonia is a polar molecule with N carrying a partial negative
charge. Molecular shape is important with respect to determining if
a molecule is polar or not.

4

Look at Valence Bond Theory (VBT)
2p

2s

N [He] 2s2 2p3

Hybridization
mix the orbitals -" like mixing together a red and white plant"
H HH

H 1s1
2

3

Hybridization of N = sp3

N [He] 2s 2p

We know that sp3 hybrids have a 109.50 angle
N
H

H
H

Molecular Structure of NH3 - cannot see the lone pair on N but
there is a flattened lone pair

N
H

H
H

The actual shape of NH3 is trigonal pyramidal (approximately tetrahedral
minus one atom).

Compared to H20
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The O in H2O has 2 bond pairs and 2 lone pairs. Two corners of the
tetrahedron are missing because they are occupied by lone pairs, not
atoms. The shape is called bent. The H-O-H angle is less than NH3, due
to the greater repulsions felt with two lone pairs

Other molecules with 2 bond plus 2 lone pairs include OF2, H2S and SF2.
Bond angles vary, but all are significantly less than 109.50.

5


The Shape of BF3


Treat this as an exception to the octet rule.
(An atom obeys the octet rule when it gains, looses or shares

electrons to give an outer shell containing eight electrons with
the configuration ns2np6). Many molecules such as neutral
compounds of Boron simply do not contain enough valence
electrons for each atom to be associated with eight electrons.

2

B 2s 2p
F 2s2 2p5

sp2

2s

2p
this leaves an empty 2p orbital
F

B

F
F
Six electrons around the Boron

6


This leaves an unused "p orbital" perpendicular to the plane of BF3
F
F

B
F

But if we want B to have an octet how can we achieve this?

A hybrid of 4 resonance structures is the best Lewis representation for the real
strucure of BF3.
F

F

B

B
F

F

F

F

F
F

B
F

F

B
F

F

However...
In this structure with a double bond the fluorine atom is
sharing extra electrons with the boron.
The fluorine would have a '+' partial charge, and the boron a
'-' partial charge, this is inconsistent with the
electronegativities of fluorine and boron.
Conclusion - the Octet Rule breaks down here.

7




Evidence for a resonance structure comes from the B-F
distances measured in the solid state. They are shorter by ~15
pm’s compared to the B-F distances in BF4-. Generally as we
move from a single bond towards a double bond our bond
lengths shorten by approximately 15 ppm’s.
F
F

B

F

F

C-C Distances CH3CH2
155 ppm

CH=CH
140 ppm

BF3 Resonance


Rehybridize the F’s to sp2

F
F

B

empty
'p' on B

F

filled
'p' on F

The MO diagram is complex but the result for BF3 is one πbond spread over 3 B-F links.
F
B

F

empty
'p' on B

F

filled
'p' on F

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To Summarize: BF3

The B atom has three bond pairs in its outer shell. Minimizing the
repulsion causes this molecule to have a trigonal planar shape, with the
F atoms forming an equilateral triangle about the B atom. The F-B-F
bond angles are all 120°, and all the atoms are in the same plane.

5.
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Lewis Acids and Bases

BF3 reacts strongly with compounds which have an unshared pair
of electrons which can be used to form a bond with the boron:

BF3 – Lewis Acid – electron pair acceptor.
NH3 – Lewis Base – electron pair donor.

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6.
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σ versus Π-bonding
Ethene, C2H4, sp2
H

H

H

H

Two lobes one with a
positive sign the other
with a negative sign go
though a node.

Nodal Plane
fn = 0 (wave function)
i.e. no electron density

p orbital not used in hybridization

The three sp2 hybrid orbitals arrange themselves as far apart as
possible - which is at 120° to each other in a plane. The remaining p
orbital is at right angles to them.
C-H overlap to give sigma bonds.

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The two carbon atoms and four
hydrogen atoms would look like
this before they joined
together:

The various atomic orbitals which are pointing towards each other
now merge to give molecular orbitals, each containing a bonding pair
of electrons.

σ orbital – no nodal planes
Π orbital one nodal plane containing the nuclei.

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Notice that the p orbitals are so close that they are overlapping
sideways.
This sideways overlap also creates a molecular orbital, but of a
different kind. In this one the electrons aren't held on the line

between the two nuclei, but above and below the plane of the
molecule. A bond formed in this way is called a pi bond.

Π-orbital above and
below nodal plane
The σ -bond is protected but the Π -bond is sticking up and is
not protected by the rest of the molecule, hence these
electrons are exposed to reacting species and it is why alkenes
and alkynes are reactive.

7. Relationship between Lewis Structure, VBT,VSEPR
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Valence Shell Electron Pair Repulsion Theory (VSEPR) enables us to
predict the shape of the central atoms electron pairs and in turn the
hybridization of the central atom.
Lewis Structure

Electron Pair Geometry (VSEPR) - non bonding electrons and bonded atoms

hybridization

molecular geometry - only looks at shape of
atoms; not lone pairs
(VBT)
bond overlap

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Electron pair geometry:

Methane,

Ammonia,

Water

109.5

107.5

104.5

Tetrahedral

Tetrahedral

Tetrahedral

Molecular geometry : Tetrahedral Triangular pyramidal Bent/Angular
Electron pair = nonbonding electrons + bonded atoms
Molecular – only looks at shape of atoms; not lone pairs

Number of Bonded Atoms
and Lone Pairs on Central
Atom

Shape (e-pair Geometry)

Hybridization

2

Linear

sp

3

Trigonal Planar/ triangular

sp2

4

Tetrahedral

sp3

5

Trigonal Bipyramidal

sp3d

6

Octahedral

sp3d2

Examples:
1. H2O

H

O

H

- electron pair geometry = tetrahedral (2 lp
and 2 bp)
- molecular shape = bent
- O hybrization = sp3

12


F

2. XeF4 (36 electrons)
F

Xe

F

F

six pairs of electrons around Xe
lone pair geometry - octahedral
Xe = sp3d2 hybridized
F = sp3 hybridized

the lone pairs are far appart therefore the compound as a SQUARE PLANAR
molecular geometry.

F

F
Xe

F

F

A typical midterm/exam question would be:
1. Draw the Lewis Structure of XeF4
2. Give (i) the molecular shape, (ii) the electron pair geometry at
the central atom and (iii) the hybridization of the central atom.

Practice Exercise


Draw the Lewis structure of BrOF3.



Give its electron pair geometry, the hybridization of the central

atom and its molecular geometry.

13


2P32 – Principles of Inorganic Chemistry

Dr. M. Pilkington

Lecture 2 - Introduction to Metal Complexes

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Metal Complexes: What are they?

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Why are so many transition metals six-coordinate?

Werner’s Coordination Theory.
Geometry of six-coordinate complexes, geometric
isomers?

Assignment 1 due in next Monday at 4.30pm, please write your name and
student ID on your work and staple all loose sheets together.

Transition Metals – located in the d-block of the periodic
table

alkali
metals

C, H, N, O, halogens
s-block

p-block
transition metals
d-block

d-block their d orbitals are filling
focus for coordination chem.

(f-block) inner transition metals

Rows are across are called Periods. Columns down are called Groups

Transition Elements begin in the 4th period
K

Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga

Ca
1

2

[Ar]4s [Ar]4s

d-block

1


Ground State Electronic Configuration for the first
row Transition Metals
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Sc - Scandium
Ti – Titanium
V-Vanadium
Cr – Chromium
Mn – Manganes e
Fe – Iron
Co – Cobalt
Ni – Nickel
Cu – Copper
Zn – Zinc

10
Ni

Pd
Pt

[Ar] 3d14s2
[Ar] 3d2 4s2
[Ar] 3d3 4s2
[Ar] 3d5 4s1
[Ar] 3d5 4s2
[Ar] 3d6 4s2
[Ar] 3d7 4s2
[Ar] 3d8 4s2
[Ar] 3d10 4s1
[Ar] 3d10 4s2

11
Cu
Ag
Au

12
Zn
Cd
Hg

Pd – Palladium
Pt – Platinum
Ag – Silver (Argentum)
Au – Gold (Aurum)
Cd – Cadmium
Hg – Mercury (hydrargyram) “liquid silver”

ƒ Most transition metals have several oxidation states.
ƒ Mn exists in 11 oxidation states -3 upto +7

ƒ Transition metals and their compounds are usually brightly
colored.

2


“Metal Complexes” (coordination compounds)
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Origin of the name originates from the 1800’s, it was known
metal had “valencies” (oxidation numbers) that could be
satisfied by combination with elements having opposite
valencies.
Cr3+ valence of +3
O2- valence of -2
Cl- valence of -1

Examples of metal complexes CrCl3, Cr2O3
BUT CrCl3 reacts with ammonia (NH3) to form a new compound.
CrCl3 + 6NH3 → CrCl3.6NH3
Called a complex because nobody at the time
understood how they formed.

Alfred Werner – late 1800’s the father of coordination
chemistry.
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Studied in Switzerland at the University of Zurich.

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He lectured in both organic and inorganic chemistry.

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He developed the theory of coordination chemistry.

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He prepared and studied coordination compounds and
discovered optically active forms of 6-coordinate
octahedral complexes.

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His coordination chemistry extended through a whole
range of systematic inorganic chemistry and into
organic chemistry and he was awarded the Nobel
Prize in Chemistry in 1913.

Nobel Lecture
/>
3

Werner studied the following metal complexes:
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CoCl3 forms four different compounds with NH3.

CoCl3.6NH3 AgNO3 (excess)
(1:6 mol ratio)
Yellow
CoCl3.5NH3
Lavender

AgNO3 (excess)

3AgCl (ppt)

2AgCl (ppt)

One Cl does not react
CoCl3.4NH3
Green
CoCl3.4NH3
Violet

AgNO3 (excess)

1AgCl (ppt)

AgNO3 (excess)

1AgCl (ppt)

Werner’s Conclusions:
1.

2.

3.

4.

5.

In this series of compounds, cobalt has a constant
coordination number of 6 (coordination number is the
number of groups that can bond directly to the metal).
As the NH3 molecules are removed they are replaced by
Cl- which acts as if it is covalently bonded to cobalt.
Chloride and Ammonia are now called ligands.
Ligands are a Lewis base/electron pair donors that can
bind to a metal ion.
A metal complex – metal ion combined with ligands.

4


6.

Coordination complexes are neutral and counter ions are not

bonded to the central metal ion but balance the charge.

For example:

+3

0

-3

[Co(NH3)6]Cl3
counter ions
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„

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The ligands directly coordinated to the metal are contained
within the square bracket.
Six NH3 bonded to Co.
3 chloride ions are not bonded to the Co these are counter ions,
they balance the charge (Co3+) they are “free” to react with
AgNO3 to give 3 moles of AgCl.

React with 3 moles of AgNO3

H2O

[Co(NH3)6]Cl3 →

Yellow

[Co(NH3)6]3+ + 3Cl-

React with 2 moles of AgNO3

[Co(NH3)5Cl]Cl2 rewrite as [Co(NH3)5Cl]2+ + 2ClLavender (now only two reactive Cl-’s).

React with 1 mole of AgNO3

[Co(NH3)4Cl2]Cl rewrite as [Co(NH3)4Cl2]+ + ClTwo isomers green and violet

Isomers - have the same formula but different structures, i.e.
different spatial arrangements

5


What is the geometry of [Co(NH3)4Cl2]Cl ?
„

Consider Hexagonal Planar
Cl

Cl
Cl

H3N

H3N

Co

Co

H3N

NH3

Cl
NH3

Co

Cl

H3N

NH3

H3N

NH3

NH3

1,2 isomer
ortho

1,3 isomer

meta

H3N

NH3
Cl
1,4 isomer
para

There are three possibilities so this does not fit with
the observations

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Consider Trigonal Prism
Cl

Cl

Co

Cl

Cl
or

Co

or

Co

Cl
Cl
There are three possibilities so this does not fit with
the observations

6


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Consider Octahedral
/>
6 vertices, 8 sides

M

Metal ion in the center, ligands are on the vertices, all six
vertices are identical.

Geometry of [Co(NH3)4Cl2]Cl is Octahedral
Cl

Co

Cl

0

180

or

900
Cl

Co

Cl

trans isomer the two Cl
ligands are far apart (1800)

cis isomer - the two Cl
ligands are close to each other (900)

Many other examples of complexes of this coordination geometry
known.
This geometry reduces the steric crowding that is a problem in
other geometries and makes them unfavourable.
We accept Werner’s conclusions, today further evidence to
confirm his conclusions is provided by X-ray crystallography.

7


Why is the coordination number 6 so common?
We have to consider the sizes (ionic radii) of the metal ions and
ligands.

Plot the ionic radii of transition metal ions, most of them are in
the range 75-90 pm which can accommodate 6 ligands and hence
favours 6-coordination.

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„

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Consider the size of common ligands such as O as in H2O, N as
in NH3 or S. For example:
O2-

126 pm

126 pm

126 pm = radius of O (1 pm = 10-12 m)

Radius Ratio Rules

The structures of many crystals can be rationalised to a first approximation
by considering the relative sizes and numbers of ions present. The radius
ratio r+/r- can be used to make a first guess at the likely coordination
number and geometry around the cation using a set of simple rules:
Value of r+/r-

Predicted
Coordination

Number of Cation

Predicted Coordination
Geometry of
Cation

< 0.15

2

Linear

0.15-0.22

3

Trigonal Planar

0.22-0.41

4

Tetrahedral

0.41-0.73

6

Octahedral

> 0.73

8

Cubic

8


To Summarize:
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For [M(H20)6]n+ in order for the metal ion to accommodate six ligands it
must have at least a 52 pm radius.
Six waters fit exactly around a metal ion with 52 pm radius (not
crowded).
At 92 pm’s the six waters have moved far enough apart that more
waters can fit.
If the metal ion is smaller than 52pm you will fit fewer H2O’s, if it is
larger then you have more space and can go to a larger coordination

number.
Hence transition metal ions display a number of coordination numbers
but 6 is very common.
N and O donors are the most abundant in biology (amino acids), S is also
present but it is allot larger than O and as a consequence Fe fits 6 O’s
but only 4 S’s. This is apparent in thiocluster compounds of Fe.

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2P32 – Principles of Inorganic Chemistry

Dr. M. Pilkington

Lecture 3 - Classification and Nomenclature
1.

Ligand Classification:
 Coordination Chemistry and Ligands
 Monodentate Ligands
 Ambidentate Ligands
 Bridging Ligands – Biological Applications
 Multi/Polydentate Chelating Ligands

2.

Naming Metal Complexes

Rodgers Chapter 2.

Theory of Coordination Chemistry
Alfred Werner (1866-1919)
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1893, age 26: coordination theory

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Nobel prize for Chemistry, 1913



Addition of 6 mol NH3 to CoCl3(aq)

Coordination compound/complex.

N
H

3+

NH3

H
M

H

N forms a coordinate
covalent bond to the

metal

H3N
H3N

Co

NH3
NH3

3Cl–

(counter ion)

NH3
ligand (coordination sphere)

1


1. Ligand Classification
Ligand – Lewis base (electron pair donor) that is bonded
to a metal ion. Ligands are anionic or neutral.
H

]3+

i.e. [Fe(NH3)6
Metal Complex
Fe

N

H

H

Ligand - Lewis Base
Has one pair of electrons

:NH3

The NH3 shares its electron with the Fe(III) metal ion.

Classes of Ligands
1. Monodentate Ligands
“one toothed” – bind to a metal ion through a single
donor site.


For example :NH3 is a monodentate ligand.

Co3+

:NH3

Fe3+

Cl

Fe2+

C

N

2


2. Bridging Ligands
Bind to two or more metal ions simultaneously.


For Example:

O2Cl
Fe

C

N

2

(H2O)5Fe
Co3+

O
Cl

neutral ligand
Fe(H2O)5
Co3+

n+

5+

Fe

[(NC)5Fe(III)CNFe(III)(CN)5]5-

Fe can exist in number of oxidation states.
A biologically important metal ion.

A Biological Application of Fe
Iron-sulfur proteins are proteins characterized by the presence
of iron-sulfur clusters containing sulfide-linked iron centers in
variable oxidation states.
Structural motifs
 In almost all Fe-S proteins, the Fe centers is tetrahedral and
the thiolato sulfur centers, from cysteinyl residues, are terminal
ligands. The sulfide groups are either two- or threecoordinated. A common motif features a four iron ions and four
sulfide ions placed at the vertices of a cubane-type structure.
4Fe-4S clusters

3