1. Section 5.4—Polarity of Molecules

2. Electronegativity
The pull an atom has for the electrons it shares with another atom in a bond.
Electronegativity is a periodic trend
As atomic radius increases and number of electron shells increases, the nucleus of an atom has less of a pull on its outermost electrons

3. Periodic Table with Electronegativies
increases
decreases

4. Polar Bond
A polar covalent bond is when there is a partial separation of charge
One atom pulls the electrons closer to itself and has a partial negative charge.
The atom that has the electrons farther away has a partial positive charge

5. Two atoms sharing equally
Each nitrogen atom has an electronegativity of 3.0
They pull evenly on the shared electrons
The electrons are not closer to one or the other of the atoms
This is a non-polar covalent bond

6. Atoms sharing almost equallyC H
Electronegativities: H = C = 2.5
The carbon pulls on the electrons slightly more, pulling them slightly towards the carbon
Put the difference isn’t enough to create a polar bond
This is a non-polar covalent bond

7. C O H Sharing unevenly Electronegativities: H = 2.1 C = 2.5 O = 3.5
The carbon-hydrogen difference isn’t great enough to create partial charges
But the oxygen atoms pulls significantly harder on the electrons than the carbon does. This does create a polar covalent bond
This is a polar covalent bond

8. Showing Partial Charges
There are two ways to show the partial separation of charges
Use of “” for “partial”
Use of an arrow pointing towards the partial negative atom with a “plus” tail at the partial positive atom C O H + - C O H

9. Ionic Bonds
Ionic bonds occur when the electronegativies of two atoms are so different that they can’t even share unevenly…one atom just takes them from the other

10. How to determine bond type
Find the electronegativies of the two atoms in the bond
Find the absolute value of the difference of their values
If the difference is 0.4 or less, it’s a non-polar covalent bond
If the difference is greater than 0.4 but less than 1.4, it’s a polar covalent bond
If the difference is greater than 1.4, it’s an ionic bond

11. If the bond is polar, draw the polarity arrow
Let’s Practice
C – H
O—Cl
F—F
C—Cl
Example:
If the bond is polar, draw the polarity arrow

12. If the bond is polar, draw the polarity arrow
Let’s Practice
C – H
O—Cl
F—F
C—Cl
2.5 – 2.1 = non-polar
3.5 – 3.0 = polar
4.0 – 4.0 = non-polar
2.5 – 3.0 = polar
Example:
If the bond is polar, draw the polarity arrow

13. Polar Bonds versus Polar Molecules
Not every molecule with a polar bond is polar itself
If the polar bonds cancel out then the molecule is overall non-polar.
The polar bonds cancel out.
No net dipole
The polar bonds do not cancel out.
Net dipole

14. The Importance of VSEPR
You must think about a molecule in 3-D (according to VSEPR theory) to determine if it is polar or not!
Water drawn this way shows all the polar bonds canceling out. O
H H
But water drawn in the correct VSEPR structure, bent, shows the polar bonds don’t cancel out!
Net dipole
H O H

15. Let’s Practice
Example:
Is NH3 a polar molecule?

16. Let’s Practice N H Example: Is NH3 a polar molecule?
Electronegativities:
N = 3.0
H = 2.1
Difference = Polar bonds
VSEPR shape = Trigonal pyramidal
Yes, NH3 is polar
Net dipole