Download - Notes-1 Periodic Trends and Lewis Structures
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CH 101 Inorganic Chemistry
1. The periodic table of elements.
2. Shapes of inorganic compounds
3. Chemistry of materials
4. Coordination compounds: Ligands, Nomenclature,
Isomerism, stereochemistry, VB,
CF & MO Theories
5. Bioinorganic chemistry
6. Organometallic chemistry.
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TEXT BOOKS for the COURSE
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Attendance : 75% must
Students with not adequate attendance will be awarded F grade
Quiz : 2 (one before midsem and one after)
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Periodic Table
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Atomic Radii Trends (in pico meters)
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First Ionization Energy (KJ/mol)
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Goldschmidt Classification
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Metallic Character Trends
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Periodic Characteristics of Compounds
Bond Enthalpies of E-H for p block elements decrease down the group where as in d block they increase.
For atoms which has no lone pair E-X bond enthalpy decreases down the group.
For atoms which has lone pairs, E-X bond enthalpy increases between Periods 2 & 3, and then decreases down the group.
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Classification of Binary Hydrides
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Acidity of Chlorides
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Why elemental sulphur forms rings or chains with S-S single bonds, where as oxygen forms diatomic molecules O=O.
Oxygen is much smaller than sulphur.
lone-pair repulsion weakens O-O single bonds, catenated
oxygen compounds very unstable. The S-S single bond is
quite strong (266 kJ mol-1) => increased catenation on
going to sulphur.
- O2(g) vs. S8(s)
-H2O2 is unstable and strongly oxidizing but H2Sn with n up to 100
-O-O bonds used in biological oxidations but S-S bonds stabilize
protein folding.
- ozonides O3 - very unstable but many polysulphides SnX.
- S-S bonds in dithionite and polythionites.
Catenation
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Concept of Bond Energies and Factors Influencing Them
� Better energy match means stronger, shorter bonds.
� Overlap, fractional population of bonding electrons under the
influence of bonding nucleus; better overlap, stronger bond.
� Core electron repulsion (important for O, F, 2nd row down).
� Non-bond electron repulsion (important at small bond distances).
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Homopolar Single Bonds
� Absolute bond strength values are not important, more important to
establish trends and reasons.
� H-H strongest; no core electron repulsion.
� Down group IV(14), bond strength goes down as overlap and core
repulsion goes up.
� C-C > N-N > O-O, F-F
� F-F << Cl-Cl non-bonded electron repulsion important.
� 2nd row onwards, across period, irregular, complex reasons.
Homopolar Multiple Bonds
� 1st row, multiple bonds strong, good overlap
� However, note N=N < O=O due to non-bonded electron repulsion
� Down group, poorer overlap because radius is higher and core electron
repulsion increases if radius drops to improve overlap
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Heteropolar Single Bonds
Element-Hydrogen Bonds
� Across period, B(E-H) increases, no core electron repulsion, so overlap
can increase as covalent radius can be small
� Down group, B(E-H) decreases, as overlap decreases, as energy match
condition is worse
Element-Fluorine Bonds
� Across period decrease, core electron repulsion increases as atoms
become smaller
� Down a group, should decrease as core electron repulsion increase and
overlap decreases: C-F > N-F > O-F > F-F order due to non-bonded
electron repulsion
� C-F vs. Si-F dπ-pπ bonding in Si-F
� N-F vs. P-F dπ-pπ and reduced non- bonded electron repulsion
� O-F vs. S-F dπ-pπ bonding
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Element-Halogen Single Bonds
The trend is always E-F > E-Cl > E-Br > E-I (true for nearly all elements), except:
F-F and Cl-F due to large non-bonded electron repulsion in F-F bond
Heteropolar Multiple Bonds
� Trends similar to homopolar bond energies
� Down group, Bond energy decrease as overlap decreases
� Across period, Bond energy decreases as core repulsion increases
� (E=O) > (E=S), overlap and core repulsion
� (E=O) > (E=N), due to overlap considerations
Changes with Oxidation State
For all cases B(E-Hal) decrease as oxidation state increases.
� Contracted central atom orbitals, overlap improves but core electron
repulsion increases.
� B(E-F) decreases least as oxidation state increases, therefore highest
oxidation state most likely with F as ligands.
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Comparison with d and f Blocks
Bond strengths generally decreases down p-block Groups, but the opposite
trend with d-block groups, with 3rd row TM’s tending to form the strongest
bonds.
The higher Oxidation States tend to be more stable for the 3rd row TM’s as well,
whereas down the p-block Groups high Oxidation States become less stable.
This in fact is the reason for the bond strength trends, as the relativistic
stabilisation and contraction of the s and p orbitals for the third row elements
leaves the d-orbitals more exposed and so overlap is better and bonds are
stronger.
A similar situation is found for the f-block elements where covalent bonding is
much more prevalent in the chemistry of the actinides.
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Anamolies & Similarities
Chemical properties of first member of each group are significanlydifferent from its congeners.
Small atomic radius, high ionization energies, high electronegativities& low coordination number
For main group elements there are similarities between atomic number Z and Z+8
For d-block elements there are similarities between atomic number Z and Z+22
Z 14 15 16 17Si P S Cl
Z+8 22 23 24 25Ti V Cr Mn
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Element Atomic radius Oxidation Potential
Al 143 pm -1.66 V (Al3+/Al)
Sc 160 pm -1.88 V (Sc3+/Sc)
ClO4– vs MnO4
– oxidizing properties
XeO4 vs OsO4 Structural Similarities
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Learning goals:1. Writing valid Lewis structures for molecular
substances
2. Predicting molecular geometry from Lewis structures (VSEPR theory)
3. Understanding electronegativity and how this concept allows the distinction between polar bonds and non-polar bonds
4. Using Lewis structures to determine whether a molecule has a dipole moment or not
5. Using the octet rule to compute formal charges on atoms and multiple bonding between atoms
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Some issues about Lewis Structures (1) Drawing “valid” Lewis structures which follow the “octet”
rule (holds almost without exception for first full row)
(2) Drawing structures with single, double and triple bonds
(3) Dealing with isomers (same composition, different
constitution)
(4) Dealing with resonance structures (same constitution,
different bonding between atoms)
(5) Dealing with “formal” charges on atoms in Lewis structures
(6) Dealing with violations of the octet rule:
Molecules which possess an odd number of electrons
Molecules which are electron deficient
Molecules which are capable of making more than four
covalent bonds
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• In 1916 G. N. Lewis proposed that atoms
combine in order to achieve a more stable
electron configuration.
• Maximum stability results when an atom
is isoelectronic with a noble gas.
• An electron pair that is shared between
two atoms constitutes a covalent bond.
The Lewis Model of Chemical Bonding
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• The order in which the atoms of a molecule are connected is called its constitution or connectivity.
• The constitution of a molecule must be determined in order to write a Lewis structure.
Constitution
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Resonance: between Lewis structures lowers the calculated energy of the
molecule and distributes the bonding character of electrons over the
molecule
Formal Charge: is the charge of an atom would have if the electron pairs
were shared equally. Lewis structure with low formal charges typically
have lowest energy
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Resonance structures – the nitrite anion: (NO2-)
ON
O ON
ON
O O
:: :
::
..:
::
:
:..
=- - -
In drawing up a Lewis dot diagram, if we are dealing withan anion, we must put in an extra electron for each negative charge on the anion:
O
N
O
:: :
::
.
:: :. -
negative charge
on anion
One extra electronin Lewis dotdiagram becauseof single negative charge on anion
Two resonance structures average structure
Bond order= 1½
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The nitrate anion:
ON
O
:: :
::
..O
....
O
NO
::
:
::
..O..
..O
N
O
::
:
: :
..O..
..
ON
O
O
- - -
-average bondorder (B.O.)=
2 + 1 + 1 = 1⅓3
B.O. = 2 B.O. = 1 B.O. = 1
to work out bond order,pick the same bond ineach structure and average the bond orderfor that bond
Number of canonical structures
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Exceptions to the octet rule: free radicals
There are some molecules that do not obey the octet rule because they have an odd number of electrons. Such molecules are very reactive, because they do not achieve an inert gas structure, and are known as free radicals. Examples of free radicals are chlorine dioxide, nitric oxide, nitrogen dioxide, and the
superoxide radical:
nitric oxide chlorine dioxide
odd electrons
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Exceptions to the octet rule.
BF3. This can be written as F2B=F with three
resonance structures. To complete its octet, BF3
readily reacts with e.g. H2O to form BF3.H2O. The actual structure of BF3 appears not to involve a double bond and does not obey the octet rule:
Possible resonancestructure for BF3, but is not importantas this wouldinvolve the very electronegativeF donating e’s to B
Best repre-sentation ofBF3 with Bhaving only6 electronsin its valenceshell
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Exceptions to the Octet rule: Heavier atoms (P, As,
S, Se, Cl, Br, I) may attain more than an octet of
electrons:
Example: PF5.
In PF5, the P atom has ten electrons in its valence shell, which occurs commonly for heavier non-metal atoms:
F
F
F
F
F
P
PF5
P has10 valenceelectrons
leave off Felectrons notshared with P
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Many phosphorus compounds do obey the
octet rule:
PF3 and [PO4]3- :
three blue electrons are
from charge on anion
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Some compounds greatly exceed an octet of electrons:
IF7 XeF6
(both I and Xe have 14 valence e’s)(Think about [XeF8]
2-)
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The same atomic composition can correspond to many Lewis
acceptable structures
Example: C6H6
This is the atomic composition of the famous organic molecule, benzene
C
CC
C
CC
H
H
H
H
H
H
How many other isomers (acceptable Lewis structures) of C6H6 are
possible?
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Isomers of the composition C6H6
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Main Implications of Lewis Structures:
Oxidation State
Hyper Valence (SO42- vs SF6)
Can Predict Bond Length, Bond Strength
Main Limitations
Fails to Predict bond angles
Structures of molecules