notes-1 periodic trends and lewis structures

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.

TEXT BOOKS for the COURSE

Attendance : 75% must

Students with not adequate attendance will be awarded F grade

Quiz : 2 (one before midsem and one after)

Periodic Table

Atomic Radii Trends (in pico meters)

First Ionization Energy (KJ/mol)

Goldschmidt Classification

Metallic Character Trends

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.

Classification of Binary Hydrides

Acidity of Chlorides

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

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).

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

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

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.

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.

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

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

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

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

• 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

• 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

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

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

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

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

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

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

Many phosphorus compounds do obey the

octet rule:

PF3 and [PO4]3- :

three blue electrons are

from charge on anion

Some compounds greatly exceed an octet of electrons:

IF7 XeF6

(both I and Xe have 14 valence e’s)(Think about [XeF8]

2-)

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?

Isomers of the composition C6H6

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