Section 5.4—Polarity of Molecules

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. increases decreases Periodic Table with Electronegativies

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. N N 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. H H C H H 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

7. H 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. + - H C O H C O H 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

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

12. Let’s Practice 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 C – H O—Cl F—F C—Cl 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. O H H H O H 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

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

16. H N H H Let’s Practice Example: Is NH3 a polar molecule? Electronegativities: N = 3.0 H = 2.1 Difference = 0.9 Polar bonds VSEPR shape = Trigonal pyramidal Yes, NH3 is polar Net dipole