Chemical Bonding and Molecular Structure (Chapter 9) - PowerPoint PPT Presentation

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Chemical Bonding and Molecular Structure (Chapter 9)

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and Molecular Structure (Chapter 9) Ionic vs. covalent bonding Molecular orbitals and the covalent bond (Ch. 10) Valence electron Lewis dot structures – PowerPoint PPT presentation

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Title: Chemical Bonding and Molecular Structure (Chapter 9)


1
Chemical Bonding and Molecular Structure
(Chapter 9)
  • Ionic vs. covalent bonding
  • Molecular orbitals and the covalent bond (Ch.
    10)
  • Valence electron Lewis dot structures
  • octet vs. non-octet
  • resonance structures
  • formal charges
  • VSEPR - 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 IE (metal) to an element of high
    electron affinity (EA) (nonmetal)
  • Na(s) 1/2 Cl2(g) ? Na Cl-
  • ? NaCl (s)

- primarily between metals (Grps 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.
6_H2bond.mov
8
Bond Formation
  • A bond can result from a head-to-head 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).
9
Sigma Bond Formation by Orbital Overlap
Two s Atomic Orbitals (A.O.s) overlap to form an
s? (sigma) Molecular Orbital (M.O.)
10
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
11
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
12
Bond and Lone Pairs
  • Electrons are distributed as shared or BOND PAIRS
    and unshared or LONE PAIRS.

This is a LEWIS ELECTRON DOT structure.
13
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

14
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
15
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.
16
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.
17
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
18
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.
19
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.
20
Double and even triple bonds are commonly
observed for C, N, P, O, and S
21
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.
22
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
23
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.
24
Violations of the Octet Rule
  • Usually occurs with
  • Boron

elements of higher periods.
25
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.
26
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.
27
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.

28
Carbon Dioxide, CO2
At OXYGEN
At CARBON
29
Carbon Dioxide, CO2 (2)
An alternate Lewis structure is

C atom charge is 0.
AND the corresponding resonance form
30
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).

31
Boron Trifluoride, BF3
What if we form a BF double bond to satisfy the
B atom octet?
32
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

33
Thiocyanate ion, (SCN)-
Which of three possible resonance structures is
most important?
ANSWER C gt A gt B
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