1. Starter
For each ion, draw a dot-and-cross diagram and predict the shape and bond angles.
H3O+
NH2-

2. Electronegativity and polarity
L.O.:
Describe electronegativity in terms of an atom attracting bonding electrons in a covalent bond.
Explain how a permanent dipole can result in a polar bond.

3. Electron density describes the way negative charge is distributed in a molecule.
Where is the electron density in the H-O?
Which atom has the most electron-pulling power?

4. Electronegativity is a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond

5. Factors affecting electronegativity
1) Nuclear charge – the more protons, the stronger the attraction from the nucleus to the bonding pair of electrons.
2) Atomic radius – the closer the bonding electrons to the nucleus, the stronger the attraction from the nucleus to the bonding pair of electrons.
3) Shielding – the less shells of electrons shielding (repelling) the bonding electrons, the stronger the attraction from the nucleus to the bonding pair of electrons.
Elicit from students. Remind that e- in the outer layer are used to make a bond.

6. Trend down a group Electronegativity decreases Atomic radius increases
More shielding
 Less attraction between nucleus and bonding pair of electrons

7. Trend across a period Electronegativity increases
Atomic radius decreases
More nuclear charge
 Stronger attraction between nucleus and bonding pair of electrons

8. H
2.1 He Li
1.0 Be
1.5 B
2.0 C
2.5 N
3.0 O
3.5 F
4.0 Ne Na
0.9 Mg
1.2 Al Si
1.8 P S Cl Ar K
0.8 Ca Sc
1.3 Ti V
1.6 Cr Mn Fe Co Ni Cu
1.9 Zn Ga Ge As Se
2.4 Br
2.8 Kr Rb Sr Y Zr
1.4 Nb Mo Tc Ru
2.2 Rh Pd Ag Cd
1.7 In Sn Sb Te I Xe Cs
0.7 Ba La
1.1 Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

12. Bonds like this are polar
Bonds like this are polar. The greater the difference in polarity, the more polar is the covalent bond

13. A permanent dipole is a small charge difference across a bond that results from a difference in the electronegativities of the bonded atoms. A polar covalent bond has a permanent dipole.

15. Why is fluorine more electronegative than chlorine
Why is fluorine more electronegative than chlorine? Fluorine is a smaller atom and when it forms a covalent bond, the shared electrons are closer to the nucleus.
pleanary

16. Pauling electronegativity values across the Periodic Table

17. CCl4 is a non-polar molecule
CCl4 is symmetrical. The dipoles act in different directions and cancel each other out. CCl4 is a non-polar molecule with polar bonds.
Discussion

18. Range of bonding from covalent to ionic

19. Arrange the following covalent bonds in order of increasing polarity: H-O, H-F, H-N.
The order of polarity is the same as the order of electronegativity of the second atom.

20. Write + and - signs to show the polarity of the bonds in a hydrogen chloride molecule
pleanary

21. Group these molecules into polar and non-polar
Group these molecules into polar and non-polar. Draw their shapes and show dipoles: CO2 Cl2 NH3 PF3 BF3 H2O CH4

22. Electronegativity and polarity
L.O.:
Describe electronegativity in terms of an atom attracting bonding electrons in a covalent bond.
Explain how a permanent dipole can result in a polar bond.
Find exam qs

23.                               
HCl polar bond showing the unequal sharing of a cloud of electron density.
HCl polar bond showing the partial (+) change on hydrogen and the partial ( -) change on chlorine.
HCl polar bond showing the direction of the dipole with an arrow pointing toward the more negative atom. The + on the opposite end also reminds us which atom is more positive.

25. Non-polar bond
Polar bond

26. Favoured by small, highly charged +ve ions, e.g. Li+, Be2+
Electronegativity Difference 4 - + X Y X Y X- Y+ X- Y+
Pure covalent
Polar ionic
Distorted ions
Pure ionic
Polar covalent
Electrons not
equally shared
Polarisation of covalent bonds
Polarisation of ions
Favoured by small, highly charged +ve ions, e.g. Li+, Be2+

27. Covalent
bonding
Ionic
bonding

28. +ve ion gets smaller and more highly charged, so –ve polarised more
NaCl
MgCl2
AlCl3
SiCl4
Mpt
801ºC
714ºC
190ºC
-70ºC
Structure
Ionic
Polar ionic
Polar covalent
Covalent
Difference in electronegativity decreases
+ve ion gets smaller and more highly charged, so –ve polarised more

29. +ve ion gets smaller and more highly charged, so –ve polarised more
BeCl2
MgCl2
CaCl2
SrCl2
BaCl2
Mpt
401ºC
714ºC
782ºC
870ºC
963ºC
Structure
Polar covalent
Ionic
Difference in electronegativity decreases
+ve ion gets smaller and more highly charged, so –ve polarised more

30. H2O - O H H + +
Bonds: polar
Molecule: polar

31. NH3 - N H H + H + +
Bonds: polar
Molecule: polar

32. CO2 - + - O C O
Bonds: polar
Molecule: non-polar

33. CCl4 C - + Cl
Bonds: polar
Molecule: non-polar