Chemical Bonding and Molecular Structure

1
Chemical Bonding and Molecular Structure

• Ionic vs. covalent bonding
• Molecular orbitals and the covalent bond
• Valence electron Lewis dot structures
• octet vs. non-octet
• resonance structures
• formal charges
• VSEPR (Valence shell electron pair repulsion)
• - predicting shapes of molecules
• Bond properties
• polarity, bond order, bond strength

2
Chemical Bonding

• Problems and questions
• How is a molecule or polyatomic ion held
  together?
• Why are atoms distributed at strange angles?
• Why are molecules not flat?
• Can we predict the structure?
• How is structure related to chemical and physical
  properties?

3
Forms of Chemical Bonds

• There are 2 extreme forms of connecting or
  bonding atoms
• Ioniccomplete transfer of electrons from one
  atom to another
• Covalentelectrons shared between atoms

Most bonds are somewhere in between.
4
Ionic Bonds

• Ionic compounds
• - essentially complete electron transfer from an
  element of low (metal) to an element of high
  electron affinity (EA) (nonmetal)
• Na(s) 1/2 Cl2(g) ? Na Cl-
• ? NaCl (s)

- primarily between metals (Groups 1A, 2A and
transition metals) and nonmetals (esp O and
halogens)
- NON-DIRECTIONAL bonding via Coulomb
(charge) interaction
5
Covalent Bonding

• Covalent bond is the sharing of the VALENCE
• ELECTRONS of each atom in a bond

Recall Electrons are divided between core and
valence electrons. ATOM core valence Na
1s2 2s2 2p6 3s1 Ne 3s1
Br Ar 3d10 4s2 4p5 Ar 3d10 4s2 4p5
6
Valence Electrons
8A
1A
2A
3A
4A
5A
6A
7A
Number of valence electrons is equal to the Group
number.
7
Covalent Bonding

• The bond arises from the mutual attraction of 2
  nuclei for the same electrons.

A covalent bond is a balance of attractive and
repulsive forces.
8
Polar and non-polar covalent bond

• Fig.
• (a) In the nonpolar covalent bond present, there
  is a symmetrical distribution of electron
  density. (b) In the polar covalent bond present,
  electron density is displaced because of its
  electronegativity.

Dipole moment, µ e (esu) x d (angstrom)
Greater the DM greater the polarity
9
Bond Formation

• A bond can result from a head-to-head/
  end-to-end overlap of atomic orbitals on
  neighboring atoms.

This type of overlap places bonding electrons in
a MOLECULAR ORBITAL along the line between the
two atoms and forms a SIGMA BOND (s).
S-s, s-p and p-p orbitals form sigma bond.
10
Sigma Bond Formation by Orbital Overlap
Two s Atomic Orbitals (A.O.s) overlap to form an
s? (sigma) Molecular Orbital (M.O.)
11
Sigma Bond Formation by Orbital Overlap
Two s A.O.s overlap to from an s ? M.O.
Similarly, two p A.O.s can overlap end-on to
from a p? M.O.
e.g. F2
12
Electron Distribution in Molecules

• Electron distribution is depicted with Lewis
  electron dot structures
• Electrons are distributed as
• shared or BOND PAIRS and
• unshared or LONE PAIRS.

G. N. Lewis 1875 - 1946
13
Bond and Lone Pairs

• Electrons are distributed as shared or BOND PAIRS
  and unshared or LONE PAIRS.

This is a LEWIS ELECTRON DOT structure.
14
Rules of Lewis Structures

• No. of valence electrons of an atom Group
  number

• For Groups 1A-4A (Li - C),
• no. of BOND PAIRS group number

• For Groups 5A-7A (N - F),
• no. of BOND PAIRS 8 - group No.

• Except for H
• (and atoms of 3rd and higher periods),
• Bond Pairs Lone Pairs 4

15
Building a Dot Structure

• Ammonia, NH3

1. Decide on the central atom never H.
Central atom is atom of lowest affinity for
electrons. In ammonia, N is central
2. Count valence electrons H 1 and N
5 Total (3 x 1) 5 8 electrons
or
4 pairs
16
Building a Dot Structure
3. Form a sigma bond between the central
atom and surrounding atoms.
4. Remaining electrons form LONE PAIRS to
complete octet as needed.
3 BOND PAIRS and 1 LONE PAIR.
Note that N has a share in 4 pairs (8 electrons),
while each H shares 1 pair.
17
Sulfite ion, SO32-
Step 1. Central atom S
Step 2. Count valence electrons S 6 3
x O 3 x 6 18 Negative charge
2 TOTAL 6 18 2 26 e- or 13
pairs

• Step 3. Form sigma bonds

10 pairs of electrons are left.
18
Thanks to all
19
Sulfite ion, SO32- (2)
Remaining pairs become lone pairs, first on
outside atoms then on central atom.

• Each atom is surrounded by an octet of electrons.

NOTE - must add formal charges (O-, S) for
complete dot diagram
20
Carbon Dioxide, CO2

• 1. Central atom __C____
• 2. Valence electrons _16_ or _8_ pairs
• 3. Form sigma bonds.

This leaves __6__ pairs. 4. Place lone pairs
on outer atoms.
21
Carbon Dioxide, CO2 (2)

• 4. Place lone pairs on outer atoms.

5. To give C an octet, form DOUBLE BONDS
between C and O.
The second bonding pair forms a pi (p) bond.
22
Double and even triple bonds are commonly
observed for C, N, P, O, and S
23
Sulfur Dioxide, SO2

• 1. Central atom S
• 2. Valence electrons 6 26 18 electrons
• or 9 pairs

3. Form pi (?) bond so that S has an octet
note that there are two ways of doing this.
24
Sulfur Dioxide, SO2
Equivalent structures called
RESONANCE STRUCTURES
The proper Lewis structure is a HYBRID of the
two.
A BETTER representation of SO2 is made by
forming 2 double bonds
Each atom has - OCTET - formal charge 0
O S O
25
Urea (NH2)2CO

• 1. Number of valence electrons 24 e-
• 2. Draw sigma bonds.

Leaves 24 - 14 10 e- pairs.
3. Complete C atom octet with double bond.
4. Place remaining electron pairs on oxygen
and nitrogen atoms.
26
Violations of the Octet Rule

• Usually occurs with
• Boron

elements of higher periods.
27
Boron Trifluoride

• Central atom B
• Valence electrons 3 37 24
• or electron pairs 12
• Assemble dot structure

The B atom has a share in only 6 electrons (or 3
pairs). B atom in many molecules is electron
deficient.
28
Sulfur Tetrafluoride, SF4

• Central atom S
• Valence electrons 6 47 34 e-
• or 17 pairs.
• Form sigma bonds and distribute electron pairs.

5 pairs around the S atom. A common occurrence
outside the 2nd period.
29
Explanation of the failure of octet rule
Sidgwicks rule of maximum covalency

• It is not essential to have 8 electron
  surrounding in an atom
• Maximum covalency depends on its period

period Maximum shared electron Atom
n1 2 H
n2 4 Li, F
n3 6 Na, Mg
n4 8 k
30
Sugdens view of singlet linkages

• Octet rule never violated
• In PCl5, 3 Cl atom linked by normal covalent bond
  and
• others 2 by sharing only 1 electron called
  singlet linkage

31
Formal Charges

• Formal charge is the charge calculated for an
  atom in a Lewis structure on the basis of an
  equal sharing of bonded electron pairs.

32
Nitric acid
Formal charge of H
..

• We will calculate the formal charge for each atom
  in this Lewis structure.

33
Nitric acid
Formal charge of H
..

• Hydrogen shares 2 electrons with oxygen.
• Assign 1 electron to H and 1 to O.
• A neutral hydrogen atom has 1 electron.
• Therefore, the formal charge of H in nitric acid
  is 0.

34
Nitric acid
Formal charge of O
..

• Oxygen has 4 electrons in covalent bonds.
• Assign 2 of these 4 electrons to O.
• Oxygen has 2 unshared pairs. Assign all 4 of
  these electrons to O.
• Therefore, the total number of electrons assigned
  to O is 2 4 6.

35
Nitric acid
Formal charge of O
..

• Electron count of O is 6.
• A neutral oxygen has 6 electrons.
• Therefore, the formal charge of O is 0.

36
Nitric acid
Formal charge of O
..

• Electron count of O is 6 (4 electrons from
  unshared pairs half of 4 bonded electrons).
• A neutral oxygen has 6 electrons.
• Therefore, the formal charge of O is 0.

37
Nitric acid
Formal charge of O
..

• Electron count of O is 7 (6 electrons from
  unshared pairs half of 2 bonded electrons).
• A neutral oxygen has 6 electrons.
• Therefore, the formal charge of O is -1.

38
Nitric acid
Formal charge of N

..

• Electron count of N is 4 (half of 8 electrons in
  covalent bonds).
• A neutral nitrogen has 5 electrons.
• Therefore, the formal charge of N is 1.

39
Nitric acid
Formal charges


..

• A Lewis structure is not complete unless formal
  charges (if any) are shown.

40
Formal Charge
An arithmetic formula for calculating formal
charge.
Formal charge
group numberin periodic table
number ofbonds
number ofunshared electrons


41
"Electron counts" and formal charges in NH4
and BF4-
7

4
42
Resonance
43
Resonance
In chemistry, resonance or mesomerism is a way
of describing delocalized electrons within
certain molecules or polyatomic ions where the
bonding cannot be expressed by one single Lewis
formula. A molecule or ion with such delocalized
electrons is represented by several contributing
structures (also called resonance structures or
canonical forms).
44
General characteristics of resonance

• Molecules and ions with resonance (also called
  mesomerism) have the following basic
  characteristics
• They can be represented by several correct Lewis
  formulas, called "contributing structures",
  "resonance structures or "canonical forms".
  However, the real structure is not a rapid
  interconversion of contributing structures.
  Several Lewis structures are used together,
  because none of them exactly represents the
  actual structure. To represent the intermediate,
  a resonance hybrid is used instead.
• The contributing structures are not isomers. They
  differ only in the position of electrons, not in
  the position of nuclei.

45
General characteristics of resonance

• Each Lewis formula must have the same number of
  valence electrons (and thus the same total
  charge), and the same number of unpaired
  electrons, if any.
• Bonds that have different bond orders in
  different contributing structures do not have
  typical bond lengths. Measurements reveal
  intermediate bond lengths.
• The real structure has a lower total potential
  energy than each of the contributing structures
  would have. This means that it is more stable
  than each separate contributing structure would
  be.

46
Contributing structures of some ion
47
Resonance Structures of Methyl Nitrite

• same atomic positions
• differ in electron positions

more stable Lewis structure
less stable Lewis structure
48
Resonance Structures of Methyl Nitrite

• same atomic positions
• differ in electron positions

more stable Lewis structure
less stable Lewis structure
49
Why Write Resonance Structures?

• Electrons in molecules are often
  delocalizedbetween two or more atoms.
• Electrons in a single Lewis structure are
  assigned to specific atoms-a single Lewis
  structure is insufficient to show electron
  delocalization.
• Composite of resonance forms more accurately
  depicts electron distribution.

50
Example

• Ozone (O3)
• Lewis structure of ozone shows one double bond
  and one single bond

Expect one short bond and one long
bond Reality bonds are of equal length (128 pm)
51
Example

• Ozone (O3)
• Lewis structure of ozone shows one double bond
  and one single bond

Resonance
52
Thanks to All
53
Formal Atom Charges

• Atoms in molecules often bear a charge ( or -).

Formal charge Group no. - 1/2 (no. bond
electrons) - (no. of LP electrons)

• The most important dominant resonance structure
  
• of a molecule is the one with formal charges
• as close to 0 as possible.

54
Carbon Dioxide, CO2
At OXYGEN
At CARBON
55
Carbon Dioxide, CO2 (2)
An alternate Lewis structure is

C atom charge is 0.
AND the corresponding resonance form
56
Carbon Dioxide, CO2 (3)
Which is the predominant resonance structure?
OR
Answer ? Form without formal charges is BETTER -
no ve charge on O

• REALITY Partial charges calculated
• by CAChe molecular modeling
• system (on CD-ROM).

57
Boron Trifluoride, BF3
What if we form a BF double bond to satisfy the
B atom octet?
58
Boron Trifluoride, BF3 (2)

fc 7 - 2 - 4 1 Fluorine
fc 3 - 4 - 0 -1 Boron

• To have 1 charge on F, with its very high
  electron affinity is not good. -ve charges best
  placed on atoms with high EA.
• Similarly -1 charge on B is bad
• NOT important Lewis structure

59
Thiocyanate ion, (SCN)-
Which of three possible resonance structures is
most important?
ANSWER C gt A gt B