Section 5.4—Polarity of Molecules

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

Periodic Table with Electronegativies

increases

decreases

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

Two atoms sharing equally

N N

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

Atoms sharing almost equally

Electronegativities: H = 2.1 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

C HH

H

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

C OH

H

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 OH

H

+ -C OH

H

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

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

Let’s Practice

Example:If the bond

is polar, draw the polarity arrow

C – H

O—Cl

F—F

C—Cl

Let’s Practice

Example:If the bond

is polar, draw the polarity arrow

C – H

O—Cl

F—F

C—Cl

2.5 – 2.1 = 0.4 non-polar

3.5 – 3.0 = 0.5 polar

4.0 – 4.0 = 0.0 non-polar

2.5 – 3.0 = - 0.5 polar

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

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. But water drawn in

the correct VSEPR structure, bent, shows the polar bonds don’t cancel out!

Net dipole

H O H

O H H

Let’s Practice

Example:Is NH3 a

polar molecule?

Let’s Practice

Example:Is NH3 a

polar molecule?

NH H

HElectronegativities:N = 3.0H = 2.1Difference = 0.9 Polar bonds

VSEPR shape = Trigonal pyramidal

Net dipole

Yes, NH3 is polar