1. CHE2060 Topic 1: Atoms, orbitals & bonding
Atoms, Orbitals & Bonding Topics:
Very quick history of chemistry…
What is organic chemistry?
Atomic models: nuclear to quantum
All about orbitals
How orbitals fill: electron configuration
Basic bonding: valence electrons & molecular orbitals
Lewis dot structures of molecules
Electronegativity & bond polarity
Resonance: a critical concept
Orbital hybridization: key to carbon’s “flexibility”
sp3
sp2 sp
Free electron pairs & radicals
VSEPR: classifying molecular geometry & orbital hybridization
Daley & Daley, Chapter 1
Atoms, Orbitals & Bonds

2. Electronegativity and bond polarity
CHE2060 Topic 1: Atoms, orbitals & bonding
Electronegativity and bond polarity

3. Bond polarity & electronegativity
Covalent bonds can be either polar or nonpolar depending on how equally (or unequally) their e- pair bonds are shared between the 2 bonded atoms.
Electronegativity is used to determine polarity of a bond between two atoms.
“Electronegativity is the ability of an atom in a molecule to attract the molecule’s electrons toward itself.”
Higher EN = stronger attraction.
More protons create more pull on e-.
More shells move ve- further from nucleus.
Bruice (2007)

4. Calculating bond polarity
Non-polar bonds share electrons equally (or evenly) between the 2 atoms.
Polar bonds share electrons unequally.
Degree of polarity is calculated: polarity (ΔEN) = | ENA – ENB|
A dipolar charge is a partial positive or negative charge on atoms joined by a polar covalent bond.
Small delta (δ) indicates dipolar charge. Here δ are shown only for polar bonds. + -
Type of bond
ΔEN
Ionic
> 2.0
Polar covalent
0.5 – 2.0
Nonpolar covalent
< 0.5 . ..
C – F:
H |
H - C - O - H | H \
C = O /
H - C - H
Use EN values to determine the polarity of each bond.
Crossed arrows are used to show the direction in which electrons are pulled. E- move toward the arrow head, leaving behind a positive charge (the cross).

5. Exercise: non-polar, polar or ionic?
For each pair of atoms:
a) Calculate bond polarities
b) Identify each bond as ionic, polar covalent or non-polar covalent
HBr
H2O
LiI
BrCl
NH3 KF
|2.1 – 2.8| = 0.7 polar covalent
|2.1 – 3.5| = 1.4 polar covalent
|1.0 – 2.5| = 1.5 polar covalent
|2.8 – 3.0| = 0.2 non-polar covalent
|3.0 – 2.1| = 0.9 polar covalent
|0.8 – 4.0| = 3.2 ionic

6. Visualizing polarity
The polarity of bonds (or within molecules) can be visualized as electron density in electrostatic potential maps.
“Hot” colors indicate higher electron density.
EPMs also show relative atomic size.
H - I
2.1 – 2.5
H - Br
2.1 – 2.8
H - Cl
2.1 – 3.0
H - F
2.1 – 4.0 δ+ δ-
Bruice (2007)

7. Neighbors can affect your polarity!
Inductive effects occur when the polarity of bond (A) is increased if its neighboring bond (B) has a much higher polarity.
But 1-chloroethane has an atom (Cl) with a higher EN. The polar C-Cl bond polarizes the neighboring C-C bond a bit.
This new polarity is induced.
H H
| |
H - C – C – H
H H
| |
H - C – C – Cl:
Ethane’s C-C bond is non- polar. .. vs ..
Field effects are similar, but occur when influencing neighbors are far from one another in terms of bonds, but are close in 3D structure.
Field effect polarization can also be caused by a molecule’s environment. Image placing a large magnet near one end of a long non-polar molecule.
While the yellow atoms are far apart if the peel is stretched out, they are close when its 3D structure is restored. So they affect one another’s polarity
D&D p.56

8. Generally, the most EN atom carries the charge.
Formal charge shows extreme polarities
If a molecule carries a charge, we calculate formal charge to show where that charge is located; which atom carries the charge.
Start with a Lewis structure.
Formal charge = (# ve-) – (dots + sticks) Calculated for each atom.
Draw Lewis structures & calculate formal charges for these molecules:
CH4
H3O+1
CH3O-1
Generally, the most EN atom carries the charge. H |
H - C – H
H - O – H ..
H - C – O:
H = 1 – (0+1) = 0
O = 6 – (2 + 3) = +1
So O carries the charge.
H = 1 – (0+1) = 0
C = 4 – (0 + 4) = 0
O = 6 – (6 + 1) = +1
So O carries the charge.
H = 1 – (0+1) = 0
C = 4 – (0 + 4) = 0
Makes sense since there is no charge.
D&D p.56-9

9. Exercise: calculating formal charge
Draw a Lewis structure & calculate formal charges for each atom.
CH3F
CH3OH2+1
CH3+1
NH3BF3
-1 :CH3
:CH2
HNO2 H |
H - C – F:
H = 1 – (0+1) = 0
C = 4 – (0 + 4) = 0
F = 7 – (6 + 1) = 0 .. H |
H - C – O – H
| |
H H
H = 1 – (0+1) = 0
C = 4 – (0 + 4) = 0
O = 6 – (2 + 3) = +1 ..
Start by counting number of ve-
H :F:
| |
H – N – B – F:
H :F:
H = 1 – (0+1) = 0
N = 5 – (0 + 4) = +1
B = 3 – (0 + 4) = -1
F = 7 – (6 + 1) = 0
Net = zero .. H |
H – C
H = 1 – (0+1) = 0
C = 4 – (0 + 3) = +1 H |
H – C:
H = 1 – (0+1) = 0
C = 4 – (2 + 3) = -1
H – N = O |
:O: ..
H = 1 – (0+1) = 0
N = 5 – (0 + 4) = +1
Single O = 6 – (6 + 1) = -1
Double O = 6 – (4 + 2) = o
Net = zero
H = 1 – (0+1) = 0
C = 4 – (2 + 2) = 0 H |
H – C:
D&D p.59

10. Dipole moment: polarity & geometry
Dipole moments are permanent partial charges created by unsymmetrical separation of charges.
Which of these molecules has a permanent dipole moment? δ- δ+ δ- δ+ H H H Cl
C C
C C Cl Cl Cl H
Here the dipoles cancel each other producing no overall molecular polarity, or dipole moment.
H&P p.7

11. Polarity varies with atomic context
The electronegativity values of H & O are constant, and always differ by 1.4.
However, the degree of molecular polarity also depends on overall molecular charge, and to some extent, geometry & shape.
These molecules all have dipole moments.
most reactive
least reactive
Which of these molecules is the most reactive?
The least reactive?
Bruice (2007)

12. Polarities of typical organic bonds
Note that bond dipole takes molecular geometry into account along with differences in electronegativity values.
Bond
Bond dipole (D)
C - F
1.53
C - Cl
1.59
C - Br
1.48
C - I
1.29
C - N
0.22
C - O
0.85
C - H
0.35
Bond
Bond dipole (D)
H - F
1.82
H - Cl
1.08
H- Br
0.82
H - I
0.44
H - N
1.32
H - O
1.54
What trends do you see here?
Critical point to remember:
Chemical reactivity of a bond increases with its increasing polarity!
D&D p.54-5