CH105
Part II: Inorganic Chemistry
The optimist sees the glass half full.
The pessimist sees the glass half empty.
The chemist see the glass completely full,
half in the liquid state and half in the vapor state.
A proton and a neutron are walking down the street.
The proton says, "Wait, I dropped an electron help
me look for it.”
The neutron says "Are you sure?"
The proton replies ……
What is the most important rule in chemistry?
Never lick the spoon!
What did the scientist say when he found 2 isotopes of helium?
HeHe
Did you hear Oxygen went on a date with potassium?
It went OK
IUPAC Nomenclature of elements
With atomic number above 100
• Digit Name Abbreviation
• 0 nil n
• 1 un u
• 2 bi b
• 3 tri t
• 4 quad q
• 5 pent p
• 6 hex h
• 7 sept s
• 8 oct o
• 9 enn e
114
Ununquadiu
m Uuq
118Ununoctium
Uuo
Money has recently been discovered to be a
not-yet-identified super heavy element.The proposed name is: Un-obtainium.
Factors Affecting Atomic Orbital Energies
• The energies of atomic orbitals are affected by
– nuclear charge (Z) and
– shielding by other electrons
• Higher nuclear charge increases nucleus-electron
interactions and lowers sublevel energy
• Shielding by other electrons reduces the full nuclear
charge to an effective nuclear charge (Zeff).Zeff is the nuclear charge an electron actually experiences. True Love !!
• Orbital shape also affects sublevel energy.
Shielding
The energy order of orbitals for a given
quantum number depends on shielding
effects (σ), effective nuclear charge (Z*)
& penetration of orbitals
Z* = Z - σ(inner electrons !!!)
How to determine or estimate the Z*?
P. S. There may be other ways of calculating these as given in the literature. Please stick
to this procedure as far as this course is concerned.
1. All e-’s in higher principal shell contribute 0 to σ
2. Each e- in the same principal shell contribute 0.35 to σ
{If the electron resides in s or p orbital}
3. Electrons in (n-1) shell: each contribute 0.85 to σ
4. Electrons in deeper shell: each contribute 1.00 to σ
Calculate the Z* for the 2p electron:
Fluorine (Z = 9) 1s2 2s2 2p5
Z* = Z – σ
P. S.: There may be other ways of calculating these as given in the literature. Please stick
to this procedure as far as this course is concerned.
Screening constant for one of the outer electron (2p):
6 six (2s2 2p4 two 2s e- and four 2p e-) = 6 X 0.35 = 2.10
2 (1s2 two) 1s e- = 2 X 0.85 = 1.70
σ = 1.70 + 2.10 = 3.80 and Z* = 9 - 3.80 = 5.20
1. All e-’s in higher principal shell contribute 0 to σ
2. Each e- in the same principal shell contribute 0.35 to σ
3. All inner shells in (n-1) and lower contributes 1.00
How to determine or estimate the Z*?
{If the electron resides in d or f orbital}
P. S. There may be other ways of calculating these as given in the literature. Please
stick to this procedure as far as this course is concerned.
Z* increases very slowly down a group for the
“valence electron”.
Example of Valence configuration as ‘ns1’
n Z Z*
H 1 1 1.0
Li 2 3 1.3
Na 3 11 2.5
K 4 19 2.2
Rb 5 37 2.2
Cs 6 55 2.2
0
2
4
6
8
10
12
14
16
H He Li Be Be C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Z* is effective nuclear charge
A neutron walks into a bar. He asks the bartender, "How much for a beer?"
The bartender offers him a warm smile and says, "For you, no charge".
Z* increases rapidly along a periodFor example, take period two
Li Be B C N O F Ne
3 4 5 6 7 8 9 10
1.3 1.9 2.4 3.1 3.8 4.5 5.1 5.8
2s1 2s2 2p1 2p2 2p3 2p4 2p5 2p6
Z* is effective nuclear charge
Penetration of Orbitals
Orbital shape causes electrons in some orbitals to “penetrate” close to the nucleus. Penetration increases nuclear attraction and decreases shielding
Radia
l pro
bability
Penetration of Orbitals
The penetration potential of an orbital
varies as: ns > np > nd > nf
The energy of the orbitals for a given n
varies as: ns < np < nd < nf
The penetration of 2s electron through the inner core is greater than
that of a 2p electron because the latter vanishes at the nucleus.
Therefore, the 2s electrons are less shielded than the 2p electrons.
penetration of 2s e- is greater than 2p
penetration increases nuclear attraction and decreases shielding
The electrons present in f are much less influenced by
the nucleus as compared to d, those present in d much
less influenced as compared to p, than s, etc.
Influence of nucleus on electrons
Two electrons (2e-) present in the same d-orbital repel
each other more strongly than do two electrons in the
same s-orbital.
It is essential to consider all contributions to the energy
of a configuration, and just not one-electron orbital
energies.
(Hostel Room Mates)
Order of filling of orbitalsPenetration and shielding have enabled atomic orbitals to be
arranged in rough order of increasing energy.
How do you fill electrons ?
Depicting orbital occupancy
for the first 10 elements.
Two electrons (2e-) present in the same d-orbital repel
each other more strongly than do two electrons in the
same s-orbital.
Therefore, occupation of orbitals of higher energy can result
in a reduction in the repulsion between electrons (for eg.,
4s), otherwise the repulsion will be more if the lower-energy
3d orbitals were occupied.
It is essential to consider all contributions to the energy of
a configuration, and just not one-electron orbital energies.
How do you fill electrons? Justification of 4s first over 3d
Experimental data show that d-block elements are
of the form 3dn4s2, with 4s orbital fully occupied.
Sc (at. No. 21) is [Ar]3d14s2
This order is followed in most cases
- but not always! (some exceptions)
Z = 24 Cr [Ar] 3d54s1; not [Ar] 3d44s2
Z = 29 Cu [Ar] 3d104s1; not [Ar] 3d94s2
Two atomic configurations do not follow the
sequence of filling of orbitals
As atomic number increases, energy of 3d orbitals
decreases relative to both 4s and 4p
At z = 29, energy of 3d becomes much lower than 4s
Hence order of filling 3d < 4s < 4p
Filling & removal in Transition elements
• Transition series: filling order: 4s, 3d
• removal order (cation formation): 4s, 3d (not 3d, 4s)
e.g. Ti [Ar] 3d2 4s2
• Ti2+ [Ar] 3d2 (not [Ar] 4s2) Why?
Ti2+ [Ar] 3d2 4s0 (not [Ar]3d04s2) Why?
• When 2 electrons are removed, regardless of where they
come from, all atomic orbitals contract (Z* increases because of
net ionic charge and reduced shielding)
• Contraction has a small effect on 4s orbital which owes its
low energy to its deep penetration
• Contraction in d orbital causes a considerable decrease in
energy – this decrease is evidently enough to lower the energy of
3d well below 4s in the ion that results from this.
“A lion runs the fastest when he is hungry.”
“In life go straight and turn right.”
r decreasesr
incre
ases
Metallic RadiusMetallic radii of 5d- block elements are expected to
be larger than that of the 4d-elements, but found
that these are not larger. Of course these are
larger than 3d- block elements.
Lanthanide Contraction
f-orbitals have poor shielding properties;
low penetrating power.
So Zeff (Z*) increases (more significantly) from left to
right (for 5d) across the period leading to more compact
atoms.
Ionisation Energy (IE)
The minimum energy needed to remove an electron from a
gas phase atom
Depends on:
(a) Size, IE decreases as the size of the atom increases
(b) Nuclear Charge (NC), IE increases with increase in NC
(c) The type of electron Shielding effect
Reasons:
(1) Average distance of 2s electron is greater than that of 1s
(2) Penetration effect
(3) Electronic configuration
1st IE: H = 1312 KJ mol-1 Li = 520 KJ mol -1
On moving across a period
1. the atomic size decreases
2. nuclear charge increases
Thus IE increases along a period
Ionisation Energy (IE)
I would like to apologize for not adding more
jokes... but I only update them.... periodically!
Electron affinity (EA)
The amount of energy associated with the
gain of electrons
The greater the energy released in the process of
taking up the extra electron, greater is the EA
The EA of an atom measures the tightness with
which it binds an additional electron to itself.
Electron affinity (EA)On moving across a period: As the size decreases, the
force of attraction by the nucleus increases. Consequently, the
atom has a greater tendency to attract added electron, i.e.,
EA increases
Generally the EA’s of metals are low while those of non-
metals are high
Halogens have high EA. This is due to their strong tendency
to change their configuration to ns2np6
On moving down a group,
the atomic size increases and therefore, the effective
nuclear attraction decreases and thus electron affinity
decreases
The process can be exo or endothermic
Electronegativitymeasure of the tendency of an element to attract
electrons to itself (from its neighbour)
On moving down the group• Z increases but Z* almost remains constant
• number of shells (n) increases
• atomic radius increases
• force of attraction between added electron and
nucleus decreases
Therefore EN decreases down the group
Electronegativity
On moving across a period left to right
• Z and Z* increases
• number of shells remains constant atomic
radius decreases
• force of attraction between added electron
and nucleus increases
Hence EN increases along a period
Trends in three atomic properties.
Hardness and Softness[Chemical but not mechanical]
An important concept of compounds formed
Rich people's dines with richer ones !!
High IE, smaller size, low polarizability -- makes Harder
Low IE, larger size, high polarizability -- makes softer
Chemical Hardness or Softness of an atom can be correlated
with ionization energy (IE), electron affinity (EA), size and
polarizability. If the IE > EA, the EA can be ignored.
The lighter atoms of a group are chemically harder
The heavier atoms of a group are chemically softer
Happiness is state of mind !!
SCN- can bind through either S or N depending upon
the HSAB nature of the metal ion. For e.g., Si or Pt
Trends are exhibited,
By keeping the metal same and changing the anion/ligand
By keeping the anion/ligand same and changing the metal
S will prefer Pt due to Soft … Soft type interactions, since
‘S’ is soft Lewis base & ‘Pt’ is soft Lewis Acid
N will prefer Si due to Hard … Hard type interactions, since
‘N’ is hard Lewis base, & ‘Si’ is hard Lewis Acid.
Oxidation States[In atomic state they are all zero].
Alkali atoms show +1 & alkaline earth shows +2
More electronegative atoms tend to form anions and
lesser electronegative atoms tend to form cations
when combined with others
Tendency of an atom to form ions with different oxidation states
(negative or positive) would depend on solvation or hydration
or ligation and lattice formation energies of the corresponding
ions. Compare this with the IE.
Why did the noble gas cry?
Because all his friends Ar- gon.
Bonding (Interaction) types:
Covalent, Non-covalent, Ionic
Non-covalent interactions are WEAK Interactions
between (atoms, molecules, compounds)
Atoms Molecules Supramolecules Materials/Solids
Hydrogen bonding interactions
Ion –molecular interactions
Vander Waal’s interactions
What do you call a tooth in a glass of water?
A one molar solution.
The name's Bond. Ionic Bond. Taken, not shared
Van der Waal’s Interactions
Three types of Van der Waal’s interactions:
a) Dipole – Dipole Interactions
b) Dipole – Induced Dipole Interactions
c) Induced Dipole – Induced Dipole
Transient Dipole – Transient Dipole
(London Dipersion Forces)
What do dipoles say in passing?
"Have you got a moment?”
Coordination numbers
Number of neighbours in interaction with the
central ion
-- Can be primary (closely interacting and/or
bonding)
-- Can be secondary (distant than the primary but
interacting – mostly no bonding)
-- All this affects the reactivity, conductivity,
electronic and magnetic properties
Coordination GeometryThe way the nearest neighbours are arranged in space,
a variety of geometries emerge:
(Main group, Transition and Lanthanides)
Linear (2) Trigonal (3)
Tetrahedral (4) Square planar (4)
Trigonal bipyramidal (5) Square pyramid (5)
Octahedral (6) Pentagonal bipyramid (7)
Singly capped octahedron (7)
Doubly capped octahedron (8)
Capped pentagonal bipyramid (9)
Decahedron (10)
Dodecahedron (12)
Linear (2)
Trigonal plane (3)
Square planar (4)
Tetrahedral (4)
Number Geometry Polyhedron
If I could rearrange the periodic table, I'd put Uranium and Iodine together.
Square pyramid (5)
Trigonal bipyramid (5)
NumberGeometry Polyhedron
Coordination No. 5
Did you hear oxygen and magnesium got together?
OMg!
Triagonal prism
Octahedral (6)
Coordination No. 6
Singly capped
octahedron (7)
Pentagonal
bipyramidal (7)
NumberGeometry Polyhedron
Coordination No. 7
Doubly capped
octahedral (8)
Heptagonal dipyramid Tricapped triagonal prism
Coordn. No. 8 (only for information, not for the exam)
Coordn. No. 9 (only for information, not for the exam)
Pentagonal Prism Bicapped square Prism
Octadecahedron Hendecahedron
Coordn. No. 10 (only for information, not for the exam)
Coordn. No. 11 (only for information, not for the exam)
Icosahedrons Cuboctahedrons
Hexagonal prism Hexagonal antiprism
Coordn. No. 12 (only for information, not for the exam)
