1. Chemical Bonding and Molecular Structure
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
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Bonding and Structure

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?
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Bonding and Structure

3. Forms of Chemical Bonds
There are 2 extreme forms of connecting or bonding atoms:
Ionic—complete transfer of electrons from one atom to another
Covalent—electrons shared between atoms
Most bonds are somewhere in between.
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Bonding and Structure

4. - NON-DIRECTIONAL bonding via Coulomb (charge) interaction
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)
Ionic Bonds
- primarily between metals
(Grps 1A, 2A and transition metals)
and nonmetals (esp O and halogens)
- NON-DIRECTIONAL bonding
via Coulomb (charge) interaction
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Bonding and Structure

5. Covalent Bonding ELECTRONS of each atom in a bond Br Br
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] s1
Br [Ar] 3d10 4s2 4p5 [Ar] 3d10 4s2 4p5
Br Br
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Bonding and Structure

6. 8A
Valence Electrons 1A 2A 3A 4A 5A 6A 7A
Number of valence electrons is equal to the Group number.
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Bonding and Structure

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
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Bonding and Structure

8. Bond Formation H Cl MOLECULAR ORBITAL along the line between
A bond can result from a “head-to-head” overlap of atomic orbitals on neighboring atoms. H Cl •• • +
Overlap of H (1s) and Cl (2p)
This type of overlap places bonding electrons in a
MOLECULAR ORBITAL along the line between
the two atoms and forms a SIGMA BOND (s).
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Bonding and Structure

9. Sigma Bond Formation by Orbital Overlap
Two s Atomic Orbitals (A.O.s) overlap to form an s (sigma) Molecular Orbital (M.O.)
6_H2pot.mov
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Bonding and Structure

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
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Bonding and Structure

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
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Bonding and Structure

12. Bond and Lone Pairs H Cl This is a LEWIS ELECTRON DOT structure.
Electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. •• H Cl •
Unshared or
lone pair (LP)
shared or bond pair
This is a LEWIS ELECTRON DOT structure.
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Bonding and 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
This observation is called the OCTET RULE
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Bonding and Structure

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
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Bonding and Structure

15. Building a Dot Structure 3 BOND PAIRS and 1 LONE PAIR.
3. Form a sigma bond
between the central atom
and surrounding atoms.
4. Remaining electrons form
LONE PAIRS to complete
octet as needed. H •• N
3 BOND PAIRS and 1 LONE PAIR.
Note that N has a share in 4 pairs (8 electrons), while
each H shares 1 pair.
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Bonding and Structure

16. 10 pairs of electrons are left.
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 = = 26 e-
or 13 pairs
Step 3. Form sigma bonds
10 pairs of electrons are left.
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Bonding and Structure

17. Sulfite ion, SO32- (2) O S Remaining pairs become lone pairs,
first on outside atoms
then on central atom. • O S ••
Each atom is surrounded by an octet of electrons.
NOTE - must add formal charges (O-, S+) for complete dot diagram
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Bonding and Structure

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.
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Bonding and Structure

19. The second bonding pair forms a pi (p) bond.
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.
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Bonding and Structure

20. H2CO
Double and even triple bonds are commonly observed for C, N, P, O, and S
SO3
C2F4
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Bonding and Structure

21. Sulfur Dioxide, SO2 1. Central atom = S
2. Valence electrons = 6 + 2*6 = 18 electrons
or 9 pairs
3. Form pi () bond so that S has an octet
— note that there are two ways of doing this.
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Bonding and Structure

22. O = S = O Sulfur Dioxide, SO2 RESONANCE STRUCTURES
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
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Bonding and Structure

23. Urea (NH2)2CO 1. Number of valence electrons = 24 e-
2. Draw sigma bonds.
Leaves = 10 e- pairs.
3. Complete C atom octet with double bond.
4. Place remaining electron pairs on oxygen
and nitrogen atoms.
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Bonding and Structure

24. Violations of the Octet Rule
Usually occurs with:
Boron
elements of higher periods.
SF4
BF3
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Bonding and Structure

25. Boron Trifluoride Central atom = B Valence electrons = 3 + 3*7 = 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.
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Bonding and Structure

26. Sulfur Tetrafluoride, SF4
Central atom = S
Valence electrons = 6 + 4*7 = 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.
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Bonding and Structure

27. Formal Atom Charges Formal charge = Group no.
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.
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Bonding and Structure

28. Carbon Dioxide, CO2 O C 6 - ( 1 / 2 ) 4 = 4 - ( 1 / 2 ) 8 = At OXYGENO C •
• • 4 - ( 1 / 2 ) 8 =
At CARBON
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Bonding and Structure

29. O C O C + + Carbon Dioxide, CO2 (2) C atom charge is 0. +
An alternate Lewis structure is: 6 - ( 1 / 2 ) = + O C •
• •
C atom charge is 0. 6 - ( 1 / 2 ) = + O C •
• • +
AND the corresponding
resonance form
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Bonding and Structure

30. O C O C + Carbon Dioxide, CO2 (3)
Which is the predominant resonance structure? O C •
• • + O C •
• • 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).
+1.46
-0.73
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Bonding and Structure

31. Boron Trifluoride, BF3
What if we form a B—F double bond to satisfy the B atom octet?
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Bonding and Structure

32. Boron Trifluoride, BF3 (2)+ F •• B •
fc = = Fluorine
fc = = 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
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Bonding and Structure

33. Thiocyanate ion, (SCN)-
Which of three possible resonance structures
is most important?
A. S=C=N
B. S=C - N
C S-C N
-0.52
-0.32
-0.16
Calculated partial charges
ANSWER:
C > A > B
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Bonding and Structure